General Chemistry
Introduction to Chemistry
What is Chemistry?
Chemistry
Definition: The branch of science that studies the composition, structure, properties, and changes of matter, including the energy changes that accompany these processes.
Important details:
It focuses on the interactions and transformations of substances, particularly through chemical reactions involving atoms and electrons.
Chemistry is often called the central science because it connects and supports different branches of knowledge, such as biology, physics, medicine, engineering, agriculture, and environmental science.
Example: When a nurse studies how a drug interacts with the body at the molecular level, they are applying chemistry as a central science that links medicine and biology.
Why Study Chemistry?
Importance of Chemistry
Definition: Chemistry helps us understand the nature of the environment around us, since everything we see, touch, taste, and smell is made of chemicals.
Important details:
Every process in daily life involves chemical interactions — from cooking and cleaning to breathing and brain activity.
Studying chemistry gives insight into how and why things happen, such as how food changes when cooked, what makes baking soda effective, how soap cleans, and why medicines affect people differently.
Chemistry explains ordinary phenomena and helps people make informed decisions about the products they use, from food and medicine to household materials.
Learning chemistry equips a person to think critically, solve real-world problems, and make safe, informed, and responsible choices in daily life.
Example: Reading a food label to check if a preservative is safe to consume is a practical use of chemistry knowledge in daily decision-making.
What Fields of Study Use Chemistry?
Applications of Chemistry Across Fields
Definition: Chemistry is not limited to laboratories; it is applied in many everyday professions and industries.
Important details:
In medicine, doctors, nurses, dentists, pharmacists, physical therapists, and veterinarians rely on chemistry to understand how drugs work in the body and how diseases are treated.
In engineering and manufacturing, chemistry is used to design stronger materials, develop new technologies, and improve industrial processes.
Chemistry is also applied in non-scientific jobs: chefs apply chemical reactions in cooking, hairdressers use chemical processes in hair treatments, farmers use chemistry to improve crop production, and firefighters understand combustion and safety through chemical principles. Artists, plumbers, truck drivers, and many other professionals also encounter chemistry in practical ways.
Specific fields where chemistry is used to solve real-world problems include: Health (medicines and vaccines), Environment (pollution monitoring, waste treatment), Agriculture (fertilizers, pesticides, soil analysis), Energy (fuel production, batteries), Textile Industry (synthetic fibers, dyes), Food Industry (preservation, flavoring, quality control), and Forensics (analysis of evidence, DNA analysis, toxicology).
Example: A forensic chemist analyzing a blood sample from a crime scene to identify a toxic substance is applying chemistry to forensic science.
Matter
Introduction
Matter
Definition: Anything that has mass and occupies space.
Important details:
All substances and materials that surround us — such as air and water, and any other products used every day — are made up of matter.
Scientists study the properties of matter and the changes it undergoes to identify and classify substances.
Knowledge of the properties and changes of matter has practical applications, including identifying unknown substances and evaluating consumer products for performance, safety, and environmental impact.
Example: Checking whether a mystery liquid is water or rubbing alcohol by comparing their smell and how quickly they evaporate uses knowledge of the properties of matter.
Properties of Matter
Properties of Matter (Overview)
Definition: The characteristics that describe and identify a material, which may be physical or chemical.
Important details:
Materials are chosen for specific purposes based on their properties. For example, a cupboard uses strong wood, nails, and hinges for durability, a raincoat uses waterproof material, and electric cables use metal that conducts electricity.
Example: Choosing rubber for the soles of shoes because it grips the floor well and resists wear is a practical use of understanding a material's properties.
Physical Properties
Physical Property
Definition: A characteristic of a substance that can be observed or measured without changing or altering its chemical composition.
Important details:
Physical properties help describe and identify substances.
Examples: color, odor, density, melting point, boiling point, mass, volume, solubility, and electrical conductivity.
Example: Noticing that a metal spoon is shiny and cold to the touch describes its physical properties.
Intensive Properties
Definition: Physical properties of matter that do NOT depend on the amount of substance present.
Important details:
Examples: color, density, luster, melting point, and boiling point.
Example: A small chip of gold and a large gold bar both have the same density (19.3 g/mL) — density doesn't change with amount.
Extensive Properties
Definition: Physical properties of matter that DO depend on the amount of substance present.
Important details:
Examples: mass, volume, length, and size.
Example: A bucket of water has more mass and volume than a single glass of water, even though both are the same substance.
Chemical Properties
Chemical Property
Definition: A characteristic that describes a substance's ability to undergo a chemical change and form new substances.
Important details:
Unlike physical properties, chemical properties can only be observed when a substance undergoes a change in composition.
Examples: flammability or combustibility, ability to rust or corrode, reactivity with acids, bases, water, or oxygen, and tendency to decompose.
Properties of matter can be likened to the "personality traits" of a substance: physical properties tell us what a substance is, while chemical properties tell us what it can do.
From the material: Paper can burn in the presence of oxygen, producing ash, carbon dioxide, water vapor, and other substances. Iron can react with oxygen and moisture, leading to rust formation. Sulfur is a tasteless, odorless, pale-yellow nonmetallic solid (physical property) that reacts with oxygen in the air to form sulfur dioxide and sulfur trioxide (chemical property).
Example: Noticing that a piece of dry leaf catches fire easily when exposed to a flame describes its chemical property of flammability.
Clinical Application from the material: Oxygen (O2) has both physical (colorless, odorless gas) and chemical (strongly supports combustion) properties. For this reason, nurses strictly enforce safety protocols, ensuring no open flames or static hazards are near a patient receiving oxygen therapy.
Changes in Matter
Changes in Matter (Overview)
Definition: Matter can undergo changes that alter either its physical form or its chemical composition, classified as physical changes or chemical changes.
Physical Changes
Physical Change
Definition: A change that alters a substance's appearance, size, shape, or state without changing its chemical composition.
Important details:
No new substance is formed during a physical change.
Although the substance's appearance changes, its chemical identity remains the same.
Examples: melting ice, boiling water, dissolving sugar in water, and cutting or crushing materials.
Example: Shredding a piece of paper into small strips is a physical change — it's still paper, just in a different shape.
Chemical Changes
Chemical Change
Definition: A change that results in the formation of one or more new substances with properties different from those of the original substance.
Important details:
During a chemical change, the composition of matter is altered.
Common indicators: color changes, gas production, precipitate formation, release or absorption of energy, and the formation of a new odor.
Common examples: burning paper, rusting iron, spoilage of food, and cooking an egg.
From the material: When sucrose (table sugar) is heated, it melts, caramelizes, and eventually decomposes into carbon and other substances — this demonstrates both physical and chemical changes.
Example: A banana turning brown and mushy as it ripens over several days is a chemical change.
Practical Applications & Substance Identification
Using Properties to Identify Substances
Substance Identification
Definition: The process of identifying substances based on their distinct physical and chemical properties.
Important details:
Properties such as density, melting point, boiling point, and reactivity play a major role in distinguishing one substance from another.
Density can be used to distinguish between materials that appear similar but are actually different substances. Chemical reactions can also help determine the identity of an unknown substance.
The ability to identify substances is important in scientific investigations, manufacturing, environmental monitoring, healthcare, and consumer safety.
From the material (Connection to Science and Healthcare Practice): Two medications can appear identical as white powders, but their melting points serve as unique fingerprints — pure Aspirin melts at exactly 136°C, while counterfeit or contaminated batches melt at lower, irregular temperatures. A lab technician can also chemically verify true aspirin by immersing it in heated water with a few drops of iron(III) chloride (FeCl3) — a genuine reaction turns the clear solution into a deep, vivid purple.
Example: A jeweler measuring the density of a ring to check if it's genuine gold or a cheaper metal alloy is using properties to identify a substance.
Chemistry of Consumer Products
Consumer Products
Definition: Items used daily by individuals and households that contain substances performing specific functions based on their physical and chemical properties.
Important details:
Includes household cleaning products, personal care products, cosmetics, medications, and food products.
Example: A bottle of dish soap is a consumer product engineered with chemical properties that help it dissolve grease.
Household Cleaning Products
Definition: Products designed to clean efficiently and conveniently, engineered with specific chemical properties to sanitize, dissolve oils, or remove dirt.
Important details:
May include general-purpose cleaners, bathroom cleaners, disinfectants, sanitizers, stain removers, and specialty cleaners.
Example: A bathroom cleaner designed to dissolve soap scum and kill bacteria on tiles is a household cleaning product.
Personal Care Products
Definition: Substances used on external parts of the body, such as the skin, hair, lips, and teeth, formulated with precise physical properties (such as viscosity and pH) to safely clean and protect the body.
Important details:
Examples: soaps, shampoos, lotions, perfumes, and other grooming products.
Example: A shampoo formulated with a specific pH to avoid irritating the scalp is a personal care product.
Active vs. Inactive Ingredients
Definition: Consumer products contain active ingredients, which perform the product's intended function, and inactive ingredients, which support the product's stability, appearance, fragrance, or delivery.
Important details:
From the material: In many liquid bleach products, the active ingredient is sodium hypochlorite (bleaching and disinfecting agent). The primary inactive ingredient is water (dilutes the chemical to a safe handling level). Other inactive ingredients include small amounts of sodium hydroxide and sodium carbonate, which act as stabilizers.
Consumers frequently select products based solely on price or marketed features without verifying if listed components pose health or environmental risks. Unchecked ingredients may pose hidden hazards, including carcinogens, neurotoxins, allergens, or severe irritants.
Product labels provide important information about ingredients, proper use, safety precautions, and storage requirements.
Example: A person checking the ingredient list on a lotion to avoid a specific allergen is applying knowledge of active and inactive ingredients.
More Examples
Physical vs. Chemical Property — A knife being made of shiny, hard steel (physical properties) that can rust if left in the rain (chemical property) shows both types side by side.
Intensive vs. Extensive Property — Splitting a chocolate bar in half doesn't change its melting point (intensive), but it does cut its mass in half (extensive).
Physical vs. Chemical Change — Freezing juice into popsicles is a physical change; letting the same juice ferment into a fizzy, sour drink is a chemical change.
Substance Identification — A chemist comparing the boiling points of two clear liquids to determine whether one is water or ethanol is using properties to identify substances.
Active vs. Inactive Ingredients — In a bottle of hand sanitizer, ethyl alcohol is the active ingredient that kills germs, while the fragrance and thickening agents are inactive ingredients.
Scientific Measurements
Introduction
Measurement
Definition: The process of comparing a known quantity, like a measuring device, to an unknown quantity or the things or objects to be measured.
Important details:
It is the process of determining how many times a certain quantity is contained in a standard measuring device.
Chemistry deals with experimental and quantitative aspects of science, and its principles are based on carefully designed experiments carried out under controlled conditions.
Example: Using a ruler to determine how many centimeters long a pencil is involves comparing the pencil's length to the standard units on the ruler.
Measurements in Matter
Physical Quantity
Definition: A physical property that can be measured and is described by a number and a unit.
Important details:
A unit is a defined quantity used as a standard of measurement, and can be fundamental (basic units) or derived (a combination of fundamental units).
Common daily units: grams/kilograms (market goods), liters (water intake), kilometers per hour (speed limits).
Example: Recording a jug of water as "1.5 liters" gives both the number (1.5) and the unit (liters), making it a proper physical quantity.
Mass
Mass
Definition: The measure of the total quantity or amount of matter in an object.
Important details:
The unit of mass is a fundamental unit, usually expressed in grams (g) or kilograms (kg); the basic SI unit is kg.
Measured using a balance; a top-loading balance is commonly used in laboratories.
Common metric measurements: 1 kg = 1,000 g = 2.205 lbs; 1 g = 1,000 mg; 1 metric ton = 1,000 kg; 1 lb = 16 oz = 453.59 g.
Example: Weighing a bag of rice on a kitchen scale and getting a reading of 2 kg is a measurement of mass.
Length
Length
Definition: Measured by determining the distance between two distinct points.
Important details:
Measured using a ruler, meter stick, or other length-measuring devices. The basic SI unit for length is the meter.
Common metric measurements: 1 km = 1,000 m = 0.62137 mi; 1 m = 100 cm = 1,000 mm = 1,000,000 μm; 1 in = 2.54 cm.
Example: Measuring the height of a doorway with a tape measure and getting 210 cm is a measurement of length.
Volume
Volume
Definition: Measured by determining the amount of space occupied by a three-dimensional object.
Important details:
It is a derived unit, with its standard SI unit being the cubic meter (m³).
Regularly shaped solids can use a formula; irregularly shaped solids use the water displacement method.
Instruments: pipette, graduated cylinder, burette, and volumetric flask (a beaker or Erlenmeyer flask can be used if high accuracy isn't required).
Common metric measurements: 1 L = 1,000 mL = 1,000 cm³ = 1 dm³ = 1.0567 qt; 1 mL = 1 cm³; 1 gal = 4 qt = 8 pt = 3.7854 L.
Formula for the volume of an irregularly shaped solid (water displacement method):
Example: Dropping a small rock into a graduated cylinder of water and measuring how much the water level rises determines the rock's volume through water displacement.
Temperature
Temperature
Definition: A measure of the intensity of heat or how hot (or cold) a system is, regardless of its size; also defined as a measure of the average kinetic energy of the particles in a sample.
Important details:
The Celsius scale is based on the properties of water; the Kelvin scale is based on the properties of gases and is the SI unit of temperature.
The Fahrenheit scale is not used in scientific measurements.
Formulas:
Fahrenheit to Celsius:
Celsius to Fahrenheit:
Celsius to Kelvin:
Step-by-step example from the material: 15°F in °C → °C = 5/9 (15 − 32) = −9.4°C.
Example: Converting a body temperature of 37°C to Fahrenheit: °F = 9/5 (37) + 32 = 98.6°F.
Sources of Errors in Measurements
Experimental Error
Definition: The difference between a measured value and the accepted (true) value.
Important details:
Accuracy describes how close a measurement is to the accepted or true value.
Precision indicates the reliability or consistency of repeated measurements.
Example: A dart player consistently landing darts close to the bullseye is both accurate and precise; landing darts tightly clustered but far from the bullseye is precise but not accurate.
Systematic Error
Definition: Errors caused by consistent, directional errors that shift all measurements in the same direction (either consistently too high or consistently too low).
Important details:
Affects the accuracy of a measurement.
May occur due to faulty calibration of measuring instruments, poorly maintained instruments, or incorrect reading of instruments by the user.
Cannot be improved by repeating the measurement if the instrument itself is faulty or misused.
From the material: Reading the volume of a clear liquid at the wrong angle causes parallax reading error, resulting in consistently high or low readings.
Example: A weighing scale that is not zeroed properly and always adds an extra 5 grams to every measurement produces a systematic error.
Random Error
Definition: Errors due to unpredictable fluctuations in measurements due to chance.
Important details:
Affects the precision of a measurement.
May occur due to the inability to read an instrument because the reading fluctuates (e.g., electronic noise) or problems in estimating a quantity between the graduations on an instrument.
Can be minimized by repeating the measurement and taking its average, or by refining the measurement method or technique.
Example: Measuring the same table's length five times with a slightly shaky hand and getting slightly different readings each time is a random error.
Significant Figures in Calculation
Significant Figures
Definition: The digits in a measurement that are certain, plus one digit that is estimated.
Important details:
Significant figures indicate the precision of measurements. A measurement with more significant figures is generally more precise because it indicates the measuring instrument can detect smaller differences.
No measuring device can make a measurement with 100% precision.
From the material: A graduated cylinder gives 38.6 mL (3 sig figs), while a buret gives 38.56 mL (4 sig figs) — the buret measurement is more precise because it has more significant figures.
Example: A digital scale reading 12.453 g is more precise than a scale reading only 12.5 g, because it has more significant figures.
Rules in Determining Significant Figures
Rules for Significant Figures
Definition: The set of rules used to count how many digits in a number are significant.
Important details (numbered rules):
All non-zero digits are significant. Example: 245 has 3 sig figs.
Zeros between non-zero digits are significant (confined zeros). Example: 20004 has 5 sig figs; 2.0004 has 5 sig figs.
Leading zeros are NOT significant (zeros at the beginning of a number). Example: 0.0053 has 2 sig figs.
Trailing zeros with a decimal point are significant (zeros at the end of a number). Example: 1.700 has 4 sig figs; 43.0 has 3 sig figs.
Trailing zeros without a decimal point may or may not be significant. Example: 6300 can have 2, 3, or 4 sig figs depending on context.
Exact numbers have infinite significant figures. Example: 12 eggs in a dozen is exact, with no uncertainty.
Example: The measurement 0.0420 g has 3 significant figures — the leading zeros don't count, but the trailing zero after the decimal does.
Significant Figures in Addition and Subtraction
Sig Figs in Addition/Subtraction
Definition: In addition and subtraction, the final answer is rounded off so that it has the same number of decimal places as the measurement with the fewest decimal places.
Important details — Rules in rounding off numbers:
If the first digit to be dropped is 4 or less, the last digit to be retained is not changed.
If the first digit to be dropped is 5 or greater, add 1 to the last digit to be retained.
Step-by-step example from the material (addition): 14.83 g + 1.4 g + 282.425 g = 298.655 → rounded to 1 decimal place (since 1.4 g has the fewest decimal places) = 298.7 g.
Step-by-step example from the material (subtraction): 0.0372 g − 0.00487 g = 0.03233 → rounded to 4 decimal places = 0.0323 g.
Example: 5.62 cm + 3.1 cm = 8.72, rounded to 1 decimal place (since 3.1 cm has fewer decimal places) = 8.7 cm.
Significant Figures in Multiplication and Division
Sig Figs in Multiplication/Division
Definition: In multiplication and division, the number of significant figures in the final answer is the same as the measurement with the fewest significant figures.
Step-by-step example from the material (multiplication): 2.36 cm × 0.00012 cm × 4.20 cm = 1.18944 × 10⁻³ → rounded to 2 sig figs (since 0.00012 cm has the fewest) = 1.2 × 10⁻³ cm³.
Step-by-step example from the material (division): 13.87 g ÷ 1.23 mL = 11.27642276 → rounded to 3 sig figs (since 1.23 mL has the fewest) = 11.3 g/mL.
Important details:
When carrying out a series of calculations, avoid rounding off after every step — keep 2 to 3 extra digits in intermediate results, and round off only the final answer to prevent the accumulation of rounding errors.
Example: 8.2 cm × 3.10 cm = 25.42 cm² → rounded to 2 sig figs (since 8.2 cm has the fewest) = 25 cm².
Exact and Inexact Numbers
Inexact Numbers
Definition: Numbers obtained through measurement that always contain some degree of uncertainty due to the limitations of measuring instruments and the skill of the person measuring.
Example: A measured mass of 25.4 g is an inexact number because a more precise scale could reveal more decimal places.
Exact Numbers
Definition: Numbers whose values are obtained by counting objects or defined quantities, with no uncertainty.
Important details:
Expressed as whole numbers; because they have infinite precision, they do not limit the number of significant figures in a calculated answer.
From the material: "A dozen" always means exactly 12; defined conversion factors, such as 1 minute = 60 seconds, are also exact.
Example: Counting exactly 30 students in a classroom is an exact number, since there's no uncertainty in a headcount.
Density Measurement
Density
Definition: The amount of mass contained in a given volume of a substance.
Important details:
It is a physical property of a substance and is not measured directly; it is calculated from mass and volume.
Density helps compare substances, determine the identity of substances, and explain why some objects float or sink.
Any substance with a density greater than water will sink; substances with lower density will float on water.
Formula:
Where: d = density, m = mass, V = volume.
Step-by-step example from the material (Sample Problem 1): An unknown substance has a mass of 40.83 g and a volume of 9.75 mL. d = 40.83 g ÷ 9.75 mL = 4.187692308 = 4.19 g/mL (rounded to 3 sig figs, since 9.75 mL has the fewest).
Step-by-step example from the material (Sample Problem 2): Copper has a density of 8.96 g/cm³. Mass of a 32.6 cm³ piece: m = dV = (8.96 g/cm³)(32.6 cm³) = 292.096 = 292 g (rounded to 3 sig figs).
Example: A sample of liquid has a mass of 25.0 g and a volume of 20.0 mL → d = 25.0 ÷ 20.0 = 1.25 g/mL.
Dimensional Analysis
Dimensional Analysis
Definition: A method used to convert one quantity to another using conversion factors.
Important details — Three steps in solving numerical problems:
Identify what important data is given (information given) in the problem.
Identify the quantity or data to be obtained (information required).
Determine what conversion factors to use, choosing the form that puts the sought-for unit in the numerator.
Important details:
For simplicity, final answers in numerical problem-solving are rounded to two decimal places only, unless otherwise specified — but keep 2 to 3 extra digits in intermediate results in multi-step calculations.
General formula: Information REQUIRED = Information GIVEN × Conversion Factor.
Step-by-step example from the material (Sample Problem 1): A technician needs 0.214 mL of benzene (density = 0.8765 g/mL). What mass is needed? m = 0.214 mL × (0.8765 g / 1 mL) = 0.187571 = 0.19 g.
Step-by-step example from the material (Sample Problem 2): An infant weighs 3.0 kg at birth. In grams: m = 3.0 kg × (1000 g / 1 kg) = 3.00 × 10³ g. In pounds: m = 3.0 kg × (2.205 lbs / 1 kg) = 6.62 lbs.
Step-by-step example from the material (Sample Problem 3): A medication order is 1.0 mg per kg body weight, for a patient weighing 180 lbs. Step 1: convert 180 lbs to kg → kg = 180 lbs × (1 kg / 2.205 lbs) = 81.63265 kg. Step 2: mg = 81.63265 kg × (1.0 mg / 1 kg) = 81.63 mg.
Example: Converting a 5.0 km distance to meters: m = 5.0 km × (1000 m / 1 km) = 5,000 m.
More Examples
Mass vs. Weight — A nurse recording a newborn's mass as 3.2 kg on a hospital scale is measuring mass, a fundamental quantity, not just "weight" in casual conversation.
Volume by Water Displacement — Dropping an oddly shaped seashell into a graduated cylinder and noting the water rises from 50 mL to 62 mL gives a shell volume of 12 mL.
Accuracy vs. Precision — A thermometer that always reads 2 degrees too high, but gives that same reading every time, is precise but not accurate.
Significant Figures — Reporting a measured length as "12.0 cm" instead of just "12 cm" communicates that the measurement is precise to the tenths place.
Density — A metal cube with a mass of 54 g and a volume of 20 cm³ has a density of 2.7 g/cm³, which matches the known density of aluminum, helping identify the metal.
Dimensional Analysis — Converting a recipe that calls for 2 cups of flour into grams by using the conversion factor 1 cup = 120 g is an everyday use of dimensional analysis.
Electronic Structure and Periodicity
The Nuclear Atom
Atom (Overview)
Definition: The basic unit of matter, composed of smaller particles called subatomic particles.
Important details:
Knowledge of subatomic particles is important because they determine an atom's identity, charge, and mass.
The Subatomic Particles: A Quick Review
Proton
Definition: A positively charged subatomic particle found inside the nucleus.
Important details:
The number of protons determines the identity of an atom and equals the atomic number. No two elements have the same number of protons.
Example: Helium has 2 protons (atomic number 2); sodium has 11 protons (atomic number 11).
Neutron
Definition: An electrically neutral particle located in the nucleus.
Important details:
Together with protons, neutrons contribute most of the atom's mass and are collectively called nucleons.
Formula: Number of neutrons = Mass number − Atomic number.
Example: An atom with a mass number of 23 and an atomic number of 11 (sodium) has 23 − 11 = 12 neutrons.
Electron
Definition: A negatively charged particle located outside the nucleus.
Important details:
Since atoms are electrically neutral, the number of electrons equals the number of protons.
Electrons are responsible for chemical bonding behavior and many chemical properties of elements.
Example: A neutral calcium atom (20 protons) has 20 electrons surrounding its nucleus.
Summary Table of Subatomic Particles
Subatomic Particle | Symbol | Charge | Relative Mass | Location |
|---|---|---|---|---|
Proton | p+ | +1 | 1 amu | Nucleus |
Neutron | n0 | 0 | 1 amu | Nucleus |
Electron | e- | -1 | 1/1836 amu | Outside the nucleus |
Isotopes, Radioisotopes, and Radiation
Atomic Notation
Definition: A way to represent an atom showing its atomic number and mass number.
Important details:
Notation (a): ZXᴬ, where X = element symbol, Z = atomic number (protons), A = mass number (protons + neutrons).
Notation (b): element–mass number (e.g., potassium–39).
From the material: ₂₀Ca⁴⁰ tells us the atom is calcium, with an atomic number of 20 and a mass number of 40.
Example: ₈O¹⁶ represents an oxygen atom with 8 protons and a mass number of 16.
Isotopes
Definition: Atoms of the same element, meaning they have the same number of protons but a different number of neutrons, resulting in different mass numbers.
Important details:
Isotopes of the same element have the same chemical properties (same number of electrons) but different physical properties (different masses).
From the material: Carbon-12, carbon-13, and carbon-14 are isotopes of carbon.
Example: Hydrogen-1 (no neutrons) and hydrogen-2 or deuterium (1 neutron) are isotopes of hydrogen.
Radioisotopes
Definition: Isotopes with unstable nuclei that spontaneously emit radiation to become more stable.
Important details:
Radioactive isotopes emit radiation through a process called radioactive decay, where the nucleus releases excess energy as radiation.
The three most common types of radiation are alpha (α), beta (β), and gamma (γ).
Example: Uranium-238 is a radioisotope that slowly decays over billions of years.
Alpha Radiation (α)
Definition: Radiation consisting of alpha particles, each containing two protons and two neutrons, identical to a helium-4 nucleus. These particles are positively charged.
From the material: ₉₄Pu²⁴⁰ → ₉₂U²³⁶ + ₂He⁴
Example: A smoke detector using americium-241, which emits alpha particles, to detect smoke particles in the air.
Beta Radiation (β)
Definition: Radiation consisting of fast-moving electrons (β⁻) or positrons (β⁺) released from unstable atomic nuclei. Beta particles are much lighter than alpha particles and carry a single electrical charge.
From the material: ₈₈Ra²²⁸ → ₈₉U²²⁸ + ₋₁e⁰; ₁₉K⁴⁰ → ₁₈Ar⁴⁰ + ₊₁e⁰
Example: Carbon-14 undergoing beta decay is the basis for radiocarbon dating of ancient organic material.
Gamma Radiation (γ)
Definition: Radiation consisting of high-energy photons — packets of pure electromagnetic energy. Gamma rays have no mass or electrical charge and travel at the speed of light.
Important details:
Gamma rays are emitted by an unstable nucleus with excess energy, often immediately after an alpha or beta decay.
From the material: ₉₄Pu²⁴⁰* → ₉₄Pu²⁴⁰ + ₀γ⁰ (the * indicates the parent nuclide has excess energy).
Example: A doctor using a gamma-emitting radioisotope like technetium-99m to produce a medical scan image.
Comparison of the Three Types of Radiation
Alpha (α) | Beta (β⁻/β⁺) | Gamma (γ) | |
|---|---|---|---|
Symbol | 2α4 or 2He4 | -1e0 / +1e0 | 0γ0 |
Charge | +2 | -1 / +1 | 0 |
Mass | relatively heavy | lighter than alpha particles | none |
Penetrating power | low | moderate | very high |
Shielding material | Stopped by paper or skin | Stopped by thin aluminum | Requires thick lead or concrete |
Uses | Smoke detectors, radiation therapy, spacecraft power supplies | Cancer treatment, medical tracers, carbon dating, PET scan | Medical imaging, sterilization, radiotherapy |
Natural Radioisotopes
Definition: Radioisotopes that occur naturally and decay over time without human intervention, forming through natural nuclear reactions or from cosmic rays.
Important details:
Found in rocks, soil, water, and even living organisms.
Examples: Carbon-14 (beta; created in the upper atmosphere, found in all living organisms), Uranium-238 (alpha; abundant in the Earth's crust), Potassium-40 (beta and gamma; found in the human body and foods like bananas and potatoes).
Example: The potassium-40 naturally present in a banana is a small source of natural background radiation.
Human-Made Radioisotopes
Definition: Synthetic or artificial isotopes created in laboratories or nuclear reactors by altering the number of protons or neutrons in an atom, produced by bombarding atoms with neutrons or protons.
Important details:
Examples: Technetium-99m (gamma), Cobalt-60 (beta and gamma), Plutonium-238 (alpha).
Example: Technetium-99m produced in a nuclear reactor for use in hospital imaging scans.
Half-Life
Definition: The time required for half of a radionuclide sample to decay.
Important details:
Not all radioisotopes decay at the same rate — some decay within seconds, others over billions of years.
From the material: For a 100 g radioisotope, after one half-life, 50 g remains; after two half-lives, 25 g remains; after three half-lives, only 12.5 g remains.
Example: Radon-222 has a half-life of only 3.82 days, so a sample decays to half its original amount in less than a week, unlike carbon-14 which has a half-life of 5,730 years.
Applications & Effects
Radioisotope Applications in Medicine
Definition: The use of radioisotopes in medicine for diagnosis and treatment.
Important details:
For diagnosis, radioactive tracers show how organs function inside the body. From the material: Technetium-99m is commonly used to image the heart, bones, and kidneys because it emits gamma radiation that specialized cameras can detect.
For treatment, radioisotopes deliver targeted radiation therapy. From the material: Iodine-131 treats thyroid disorders such as hyperthyroidism and thyroid cancer by targeting the thyroid gland.
Diagnostic radioisotopes typically emit gamma rays or positrons (detectable without causing significant tissue damage), while therapeutic radioisotopes typically emit beta or alpha particles (which destroy diseased tissue).
Criteria for choosing medical radioisotopes: appropriate half-life, targets specific organs/tissues, suitable type of radiation, safe radiation dose, and easy to produce/obtain.
Important details — Effects in Medicine:
Positive: early disease detection, improved diagnosis accuracy, non-invasive imaging, targeted treatment of cancer, reduced need for exploratory surgery.
Negative: exposure to ionizing radiation, possible damage to healthy cells, side effects during treatment (fatigue, nausea, tissue irritation), radiation safety concerns, small increase in long-term cancer risk with excessive exposure.
Example: A patient undergoing a PET scan using fluorine-18 to detect cancer is benefiting from a medical application of radioisotopes.
Radioisotope Applications in Industry
Definition: The use of radioisotopes in industry to improve product quality, workplace safety, manufacturing efficiency, and food preservation.
Important details:
Used to detect leaks in pipelines, control material thickness (paper, plastic, metal), and perform non-destructive testing (detecting cracks in welds and metal parts using gamma rays from isotopes like cobalt-60).
Gamma rays are used for inspecting metal structures because they can penetrate thick materials; beta particles are used for measuring thickness because they are less penetrating, making changes in their absorption useful for thin materials. Alpha particles are rarely used industrially due to their extremely poor penetrating power.
Important details — Effects in Industry:
Positive: improved product quality, enhanced safety, increased manufacturing efficiency, enhanced food safety and shelf life, reduced waste and spoilage.
Negative: risk of radiation exposure to workers, potential health risks if mishandled, environmental contamination from improper disposal, high cost and complexity of safety regulations.
From the material: A common misconception is that irradiated food becomes radioactive — in reality, food irradiation kills harmful microorganisms, but the food itself never becomes radioactive, just like a person doesn't become radioactive after an X-ray.
Example: A factory using cobalt-60 to inspect welded pipeline joints for hidden cracks without cutting them open.
Radioisotope Applications in Environmental Science
Definition: The use of isotopes as tools for studying natural processes such as water movement, dating ancient materials, and monitoring pollution.
Important details:
Isotopes of hydrogen and oxygen in water help scientists understand the water cycle.
Carbon-14 is used to date ancient organic materials, helping determine the age of fossils and archaeological artifacts.
Radioactive tracers help monitor pollution by tracking contaminant movement through air, soil, and water.
Environmental scientists use both natural and artificial radioisotopes — natural radioisotopes help study natural processes and determine ages, while artificial ones can be used as tracers or contamination indicators.
Important details — Effects in Environmental Science:
Positive: monitors pollution, tracks movement of water/nutrients, determines the age of fossils/sediments, improves understanding of environmental and climate changes.
Negative: improper handling can expose researchers to radiation, accidental leaks may contaminate ecosystems, radioactive waste requires careful disposal.
Example: An archaeologist using carbon-14 dating to determine that an ancient wooden artifact is 3,000 years old.
Development of Quantum Theory
Quantum Mechanical Model
Definition: The modern model of the atom, based on probability rather than certainty, which uses mathematical equations to describe electrons as having both wave-like and particle-like properties.
Important details:
It introduces orbitals as regions where electrons are most likely to be found.
This model forms the foundation of quantum theory — the science that explains how matter and energy behave at the atomic and subatomic levels.
Earlier atomic models (Dalton's billiard ball model, Thomson's plum pudding model, Rutherford's nuclear model) were important stepping stones but could not fully explain the unusual behavior of electrons.
Example: Predicting the general region where an electron in a hydrogen atom is likely to be found, rather than its exact path, uses the probability-based approach of the quantum mechanical model.
Contribution of Max Planck
Max Planck's Contribution
Definition: Planck's 1900 proposal that energy is released in tiny packets rather than in a continuous flow, which solved the puzzle of blackbody radiation.
Important details:
Scientists at the time believed hot objects released energy smoothly and continuously, but this didn't match experimental results.
Planck's ideas became the basis of quantum theory, leading to modern technologies like smartphones, computers, lasers, and advanced medical imaging.
Example: The basic idea that light and energy come in discrete "packets" rather than a smooth stream is why we can now build devices like lasers, which rely on precise packets of energy.
Contribution of Niels Bohr
Niels Bohr's Contribution
Definition: Bohr's 1913 atomic model, in which electrons move in specific paths called energy levels or orbits around the nucleus.
Important details:
Each energy level has a fixed amount of energy; electrons move between levels by absorbing or emitting energy in the form of quanta, but cannot stay between orbits.
This model explained why atoms emit light at specific wavelengths or colors of light.
Spectral lines shown for electron drops to n = 2 (rather than n = 1) are because these transitions produce visible light; drops to n = 1 release more energy, producing invisible ultraviolet light.
Example: The colorful light seen in fireworks comes from electrons in metal atoms dropping to lower energy levels and releasing visible light, an idea explained by Bohr's model.
Contribution of Erwin Schrödinger
Erwin Schrödinger's Contribution
Definition: Schrödinger's development of wave mechanics, introducing in 1926 the Schrödinger equation, which describes how electrons behave like waves and are governed by probability.
Important details:
Instead of fixed orbits (as Bohr suggested), electrons are found in regions of space called orbitals, also known as electron clouds.
The Schrödinger equation calculates the likelihood of finding an electron in these regions, meaning an electron's path is probabilistic, not definite.
His theory showed electrons behave as both waves and particles, supporting wave-particle duality.
Example: A cloud-like diagram showing where an electron is most likely to be found around a nucleus, rather than a single fixed path, illustrates Schrödinger's electron cloud model.
Energy Levels, Sublevels, and Orbitals
Energy Levels, Sublevels, and Orbitals (Overview)
Definition: The organization of the space around the nucleus, where electrons are most likely found, divided into energy levels, sublevels, and orbitals.
Important details:
Analogy: the atom is like a building — each floor is an energy level, each room on a floor is a sublevel, and specific spots within each room are orbitals.
Each orbital can hold only two electrons, and those electrons must have opposite spins.
Main Energy Levels, n
Main Energy Level (n)
Definition: The main energy level, or shell, which describes the average distance of electrons from the nucleus.
Important details:
As n increases (n = 1, 2, 3...), electrons are farther from the nucleus, have more energy, and occupy larger regions of space that can hold more electrons.
Electrons are arranged in energy levels (K, L, M, N or 1, 2, 3, 4), starting nearest the nucleus (n = 1).
The ground state is the lowest possible energy level electrons normally occupy, and is the most stable. When energy is absorbed, electrons jump to higher levels, entering an excited state. As they return to lower levels, they release energy as light.
The outermost occupied energy level is the valence shell, and its electrons are called valence electrons.
Example: An electron in a neon sign absorbing electrical energy and jumping to an excited state, then releasing that energy as visible light when it returns to the ground state, is why neon signs glow.
Energy Sublevels and Orbitals
Orbitals and Electron Clouds
Definition: Atomic orbitals, often illustrated as electron clouds, are regions around the nucleus where electrons are most likely to be found.
Important details:
Unlike Bohr's circular paths, orbitals describe the general shape, size, and energy of the region an electron occupies, not an exact path.
Each orbital holds a maximum of two electrons, because electrons have 2 opposite spin states (+1/2 and −1/2), represented as ↑↓.
Example: Two electrons sharing the same orbital must always spin in opposite directions, similar to two magnets needing to be arranged in a stable, opposing configuration.
Energy Sublevels (Subshells)
Definition: Groups of orbitals within the same main energy level that have the same shape but different orientations in space, and the same energy.
Important details:
Labeled s, p, d, and f. The s sublevel has 1 orbital, p has 3 orbitals, d has 5 orbitals, and f has 7 orbitals.
Energy of the sublevels increases in the order s < p < d < f.
Important details:
Sublevel | Number of Orbital Types | Orbitals |
|---|---|---|
s | 1 | s |
p | 3 | px, py, pz |
d | 5 | dxy, dx²-y², dxz, dz², dyz |
f | 7 | 7 distinct f orbitals |
From the material: The first energy level (n = 1) has only the 1s sublevel and can hold a maximum of 2 electrons. The second energy level (n = 2) has 2s and 2p sublevels, holding a total of 8 electrons. The third energy level (n = 3) has 3s, 3p, and 3d sublevels, holding a total of 18 electrons.
Example: The 2p sublevel, having 3 orbitals, can hold up to 6 electrons total (2 electrons per orbital × 3 orbitals).
Summary Table of Energy Levels, Sublevels, Orbitals, and Maximum Electrons
Energy Level | Sublevel Present | Orbitals | Maximum Electrons |
|---|---|---|---|
1 | 1s | 1 | 2 |
2 | 2s, 2p | 1, 3 | 8 |
3 | 3s, 3p, 3d | 1, 3, 5 | 18 |
4 | 4s, 4p, 4d, 4f | 1, 3, 5, 7 | 32 |
Electronic Configurations
Electron Configuration
Definition: A shorthand representation that shows which energy levels, sublevels, and orbitals of an atom are filled with electrons.
Important details:
Written with three pieces of information: the number before the letter (main energy level, n), the letter (sublevel type: s, p, d, f), and the superscript (number of electrons in that sublevel).
From the material: In the configuration 2p³, the "2" means second energy level, "p" means p sublevel, and "3" means 3 electrons are located in the p orbitals of the second energy level.
Example: The configuration 3s² tells us that 2 electrons occupy the s sublevel of the third energy level.
Aufbau Diagram
Definition: A mnemonic device using diagonal arrows to show the order in which orbitals are filled, from lowest to highest energy.
Important details:
Electrons occupy orbitals in order of increasing energy, starting from the lowest, always starting with the s sublevel in each energy level since it has the lowest energy.
Example: Using the Aufbau diagram to determine that the 4s orbital fills before the 3d orbital, even though 3 is a lower main energy level number than 4.
Orbital Diagram
Definition: A pictorial representation of the electron configuration, showing individual orbitals, the distribution of electrons among them, and whether electrons are paired or unpaired.
Example: Drawing boxes for each orbital with up and down arrows to show electron pairing is an orbital diagram.
Rules in Writing the Electronic Configuration
Rule 1: Pauli's Exclusion Principle
Definition: Only two electrons can occupy the same orbital, and they must have opposite spins.
From the material: Hydrogen (1 electron) has the configuration 1s¹; helium (2 electrons) has 1s².
Example: An orbital diagram for beryllium showing both electrons in the 2s orbital with opposite spin arrows (↑↓) illustrates this rule.
Rule 2: Aufbau Principle
Definition: Electron subshells are filled in order of increasing energy (s < p < d < f); electrons occupy the first available orbital with the lowest energy before filling higher-energy orbitals.
Important details:
Within each energy level, s fills first, then p, then d, then f.
Sometimes a higher energy level orbital (like 4s) fills before a lower one (like 3d) — the diagonal rule (Aufbau diagram) serves as a guide.
From the material: Lithium's third electron enters the 2s orbital (not staying in the already-filled 1s), since 2s is the next available orbital with the lowest energy: 1s²2s¹.
Example: Boron's electron configuration, 1s²2s²2p¹, shows electrons filling the s sublevel completely before starting to fill the p sublevel.
Rule 3: Hund's Rule of Maximum Multiplicity
Definition: Electrons occupy all the orbitals of a given sublevel singly (with the same/parallel spins) before pairing begins.
Important details:
The most stable arrangement is achieved when the subshell has the greatest number of parallel spins.
From the material: Carbon's 6th electron will not pair with the electron already in the 2px orbital; instead, it occupies the next empty 2p orbital (1s²2s²2px¹2py¹). Oxygen's 8th electron finally begins pairing, since all three 2p orbitals already have one electron each (1s²2s²2px²2py¹2pz¹).
Example: Nitrogen's electron configuration, 1s²2s²2px¹2py¹2pz¹, shows each of the three 2p orbitals holding one unpaired electron before any pairing occurs, following Hund's Rule.
Noble Gas Notation
Definition: A shorthand way of writing the electron configuration of an element by substituting the last noble gas before the element to represent all inner electrons.
Important details — Guidelines:
Look for the last noble gas before the element.
Write the symbol of the noble gas (with its atomic number) enclosed in square brackets.
Write the remaining configuration starting with the s orbital of the next energy level after the noble gas.
Valence electrons are determined by counting the electrons (superscript) in the outermost energy level (highest value of n).
From the material: Carbon's noble gas notation is [2He]2s²2p², with 4 valence electrons.
Example: Sodium's noble gas notation is [10Ne]3s¹, with 1 valence electron.
The Periodic Table
Mendeleev's Development of the Periodic Table
Dmitri Mendeleev
Definition: A Russian scientist who discovered that arranging elements in order of increasing atomic mass caused certain properties to repeat regularly, and grouped elements with similar properties together.
Important details:
Mendeleev's arrangement became the foundation of the periodic table.
He left gaps in his version of the periodic table, believing some elements had not yet been discovered, and was even able to predict the properties of these unknown elements — which were later discovered and matched his predictions.
Example: This is similar to how a librarian organizing books by topic could predict there's a "missing book" if a series has a gap, even before finding the actual book.
The Modern Periodic Table
Modern Periodic Table
Definition: The current version of the periodic table, the result of the combined efforts of scientists such as Dobereiner, Newlands, Mendeleev, Meyer, and Moseley.
Important details:
Henry Moseley improved the periodic table by arranging elements in order of increasing atomic number, forming the basis of the modern periodic law: the physical and chemical properties of elements are periodic functions of their atomic numbers.
There are 118 elements in the modern periodic table. The majority are metals, very few are metalloids, and the rest are nonmetals.
The horizontal arrangement is called a period (7 periods total). The two rows below the table are the lanthanides and actinides (inner transition metals).
The vertical arrangement is called a group or family, classified into Group A (representative/main group elements) and Group B (transition elements/metals), or numbered 1 to 18 in the new classification.
Example: Chlorine and fluorine, found in the same group (VIIA), share similar chemical properties because of their identical valence electron arrangement.
The Periodic Table and The Electronic Configuration
The Period
Definition: The period number corresponds to the number of occupied energy levels in an atom, representing the valence shell (outermost energy level containing electrons).
From the material examples: Boron [He]2s²2p¹ → period 2; Potassium [Ar]4s¹ → period 4; Magnesium [Ne]3s² → period 3; Zirconium [Kr]5s²4d² → period 5.
Example: An element with the electron configuration ending in 6s² is in period 6.
The Group
Definition: Elements in the same column (group) of the periodic table have similar chemical properties because they have the same number and similar arrangement of valence electrons.
Important details:
For representative/main-group elements, the group number corresponds to the number of valence electrons in the outermost shell.
d and f electrons are always in a lower energy level and are not considered valence electrons.
Group A Number | Number of Valence Electrons |
|---|---|
IA | 1 |
IIA | 2 |
IIIA | 3 |
IVA | 4 |
VA | 5 |
VIA | 6 |
VIIA | 7 |
VIIIA | 8 |
From the material: Beryllium (Group II) [2He]2s² has 2 valence electrons; Carbon (Group IV) [2He]2s²2p² has 4 valence electrons; Bromine (Group VII) [18Ar]4s²3d¹⁰4p⁵ has 7 valence electrons.
Example: An element in Group VA with the configuration ns²np³ has 5 valence electrons.
The Blocks
Definition: Divisions of the periodic table based on the type of orbital being filled last (s, p, d, or f), which helps locate the element.
Important details:
s-block — Groups IA and IIA, where s orbitals are filled.
p-block — Groups IIIA to VIIIA, where p orbitals are filled. Group VIII-A elements have complete electron shells and are called inert/noble gases.
d-block — the transition metals (Groups 3B to 2B), where d orbitals are filled; has 10 columns for the 10 possible d electrons.
f-block — the lanthanides and actinides (28 elements total), where f orbitals are filled; has 14 columns for the 14 possible f electrons.
Group Number | Block | Last Electron Configuration |
|---|---|---|
Group IA | s-block | ns¹ |
Group IIA | s-block | ns² |
Group IIIA | p-block | ns²np¹ |
Group IVA | p-block | ns²np² |
Group VA | p-block | ns²np³ |
Group VIA | p-block | ns²np⁴ |
Group VIIA | p-block | ns²np⁵ |
Group VIIIA | p-block | ns²np⁶ |
Group B (transition metals) | d-block | ns²(n-1)d¹⁻¹⁰ |
Group B (inner transition metals) | f-block | ns²(n-2)f¹⁻¹⁴ |
From the material examples: An element with the last configuration 7s² is in period 7, s-block, group IIA — it is Radium (Ra). An element with 4s²3d³ is in period 4, d-block, group VB — it is Vanadium (V). Tellurium (Te), in period 5, p-block, has the last configuration 5s²5p⁴.
Example: An element with the last electron configuration 3s²3p⁵ belongs to the p-block, period 3, Group VIIA — it is chlorine (Cl).
Periodic Trends
Trend in Atomic Radius
Atomic Radius
Definition: The size of an atom, measured as one-half the distance between 2 atomic centers (nuclei) of two identical atoms, since an atom has no definite outer boundary.
Important details:
Across a period (left to right), atomic radius decreases. As atomic number increases, the number of protons and electrons increases, but the number of inner electrons stays the same, so the increasing nuclear charge pulls electrons closer, making the atom smaller.
Down a group (top to bottom), atomic radius increases, because more electron shells are added farther from the nucleus, and inner electrons block (shield) the nucleus from strongly attracting the outermost electrons — this is called the shielding effect.
Formula (approximate effective nuclear charge):
Where the number of inner electrons = atomic number of the noble gas before the element.
From the material: Across the third period (Na to Ar), the effective nuclear charge increases from +1 to +8, since the number of inner electrons stays at 10 while the number of protons increases from 11 to 18.
Example: Potassium (K) has a larger atomic radius than sodium (Na) because potassium is one row below sodium in the same group, meaning it has an extra electron shell.
Trend in Atomic Mass
Atomic Mass
Definition: The mass of an atom, which generally increases as one moves across a period from left to right or down a group.
Important details:
Across a period and down a group, each element has more protons and neutrons, making it heavier.
From the material: Across Period 2, lithium (6.94), beryllium (9.01), and boron (10.81) show increasing atomic mass. Down Group IA, hydrogen (1.01), lithium (6.94), and sodium (22.99) also show increasing atomic mass.
Example: Comparing fluorine (about 19.00) and chlorine (about 35.45) in the same group shows chlorine has a greater atomic mass, since it is farther down the group.
More Examples
Subatomic Particles — A neutral chlorine atom (17 protons) has 17 electrons; if it gains 1 electron to become Cl⁻, it now has 18 electrons but still only 17 protons.
Isotopes — Chlorine-35 and chlorine-37 are isotopes of chlorine, both with 17 protons but 18 and 20 neutrons respectively.
Half-Life — If a hospital receives 40 mg of a radioisotope with a half-life of 6 hours, only 5 mg remains after 18 hours (3 half-lives).
Electron Configuration — Writing out oxygen's electron configuration (1s²2s²2p⁴) shows exactly how its 8 electrons are distributed among the available sublevels.
Periodic Trends — Comparing lithium and cesium (same group), cesium has a much larger atomic radius because it has more electron shells shielding its outer electron from the nucleus.
Writing and Naming Chemical Formulas of Inorganic Compounds
Introduction
Chemical Formula
Definition: A shorthand way of writing the name of a compound, consisting of symbols that indicate the (chemical) composition of a substance.
Important details:
It tells us the types and number of atoms that are chemically combined in one unit (mole) of the substance.
A symbol represents one atom of an element, while a molecular formula represents one molecule or formula unit of a compound.
A structural formula shows the arrangement of atoms in a molecule, while a line-angle (skeletal) formula is a shorthand way of representing a molecule's bonding, where every corner and end of a line represents a hidden carbon atom (and its attached hydrogens).
Numbers written as subscripts show the number of atoms of an element combined in one unit of the molecule; if the number is only one, "1" is not written.
Example: The molecular formula C4H10 (butane) tells us the molecule is made of 4 carbon atoms and 10 hydrogen atoms.
Ionic vs. Molecular Compounds
Definition: Inorganic compounds can be classified as ionic or molecular compounds.
Important details:
Ionic compounds are made up of a cation and an anion; they can be binary (two elements) or ternary (three elements).
Molecular compounds are made up of nonmetals.
Two basic rules for writing chemical formulas: (1) the cation is written first, followed by the anion; (2) compounds are electrically neutral, so the sum of the charges of the cation and anion must be equal.
Example: Sodium chloride (NaCl) is a binary ionic compound made of a sodium cation and a chloride anion.
Rules in Writing Chemical Formulas
Steps in Writing a Chemical Formula: Ionic Compounds
Writing Ionic Compound Formulas (4 Steps)
Definition: The step-by-step process for writing the correct chemical formula of an ionic compound.
Important details (in order):
Write down the ions that combine to form the compound, cation first, then anion — the charge is always written as a superscript.
Crisscross the charges of the ions (without the signs); the interchanged numbers become subscripts of the elements in the formula. The subscript "1" is not written.
Use parentheses for polyatomic ions when there are multiple numbers of an individual polyatomic ion, to separate the polyatomic ion from the subscript.
Simplify the formula to its simplest form by dividing the subscripts by any common multiple.
From the material: Magnesium chloride: Mg²⁺ and Cl⁻ → crisscross → MgCl2. Iron(III) oxide: Fe³⁺ and O²⁻ → crisscross → Fe2O3. Ammonium sulfide: NH4⁺ and S²⁻ → (NH4)2S. Calcium oxide: Ca²⁺ and O²⁻ → crisscrosses to Ca2O2, simplified to CaO.
Example: For aluminum sulfide, Al³⁺ and S²⁻ crisscross to give Al2S3, which is already in its simplest form.
Steps in Writing a Chemical Formula: Molecular Compounds (Binary)
Greek Prefixes for Molecular Compounds
Definition: For compounds of two nonmetals, Greek prefixes tell us the number of atoms that are chemically combined.
Important details:
The prefix "mono" is not written for the cation (first element).
Common prefixes: mono (1), di (2), tri (3), tetra (4), penta (5), etc.
From the material: diphosphorus pentasulfide = P2S5; carbon dioxide = CO2; sulfur trioxide = SO3.
Example: Dinitrogen tetroxide has the formula N2O4, based on the prefixes "di" (2) and "tetra" (4).
Nomenclature of Inorganic Compounds
Trivial vs. Systematic Names
Definition: Compounds can be named using their trivial (common) name, which has no systematic relationship to the formula, or their systematic chemical name, which is directly tied to the formula.
From the material: Baking soda (NaHCO3) vs. sodium bicarbonate; muriatic acid (HCl) vs. hydrochloric acid; table salt (NaCl) vs. sodium chloride.
Example: "Lime water" is a trivial name, while "calcium hydroxide" (Ca(OH)2) is its systematic name.
Naming Cations
Definition: Cations are positive monatomic ions (except ammonium), usually metallic, and named the way their metals are named.
Important details:
Metals that form only one cation — named the same as the metal, with "ion" added. Example: Na+ is the sodium ion.
Metals that form more than one cation — can use the stock system (cation name + Roman numeral in parentheses, e.g., Cu+ = copper(I), Cu2+ = copper(II)) or the older -ous/-ic suffix system (Latin name; higher charge = -ic, lower charge = -ous, e.g., Cu+ = cuprous, Cu2+ = cupric).
Example: Iron(II) and iron(III) are named using the stock system, while ferrous and ferric are the equivalent older Latin-based names.
Naming Anions
Definition: Anions are negatively charged ions, which can be monatomic or polyatomic.
Important details:
Monatomic anions are named by adding the suffix -ide to the root of the parent element's name (root + -ide). Examples: Cl− is chloride, O2− is oxide, N3− is nitride.
Polyatomic anions (oxoanions) contain a metal or nonmetal plus one or more oxygen atoms. When only two oxoanions exist for an element, more oxygen = -ate suffix, fewer oxygen = -ite suffix. Examples: NO3− is nitrate, NO2− is nitrite.
For series with up to 4 members, per- identifies the most oxygen (ClO4− = perchlorate), and hypo- identifies the fewest oxygen (ClO− = hypochlorite).
From the material: The hypochlorite ion is the active ingredient in laundry bleach, while compounds containing the perchlorate ion can explode when they contact organic substances.
Example: SO4²⁻ is sulfate, while SO3²⁻ is sulfite, following the -ate/-ite naming rule.
Rule 1: General/Ionic Compounds
General Naming Rule
Definition: Compounds named using the cation name followed by the anion name.
Formula:
Important details:
This rule applies to bases (cation is a metal, anion is always hydroxide OH−), metallic oxides (anion is always oxide O2−), and salts (cation is not H+, anion is not OH− or O2−).
From the material examples: Na2S = sodium sulfide; CaO = calcium oxide; LiOH = lithium hydroxide; PbO2 = lead(IV) oxide (plumbic oxide); Fe(OH)3 = iron(III) hydroxide (ferric hydroxide).
Example: KMnO4 is named potassium permanganate, following this general rule.
Rule 2: Compounds of Two Nonmetals
Nonmetal Naming Rule
Definition: Compounds of two nonmetals are named using Greek prefixes to specify the number of atoms combined.
Formula:
Important details:
"Mono" is not used on the cation. If a prefix ends in "a" or "o" and the following word starts with "a" or "o," the vowel is dropped.
This rule applies to naming nonmetallic oxides.
From the material examples: SO2 = sulfur dioxide; N2O = dinitrogen monoxide; PF3 = phosphorus trifluoride; CS2 = carbon disulfide.
Example: Cl2O3 is named dichlorine trioxide, following the Greek-prefix rule.
Rule 3: Acids
Acid
Definition: A compound containing hydrogen ions (H+) as the cation and an anion, which can be a nonmetal or oxoanion.
Important details:
Acids can be binary (2 elements) or ternary (3 elements).
Binary Acids
Definition: Acids whose anion is a monatomic nonmetal (except cyanide).
Important details:
In aqueous solutions: named using hydro- + root word + -ic + "acid." Formula: hydro + root word + ic acid.
In gaseous/pure liquid form: named using "hydrogen" + the monatomic anion name. Formula: hydrogen + name of monatomic anion.
From the material: HCl(g) = hydrogen chloride; HCl(aq) = hydrochloric acid; HF(aq) = hydrofluoric acid; HF(g) = hydrogen fluoride.
Example: H2S(aq) is named hydrosulfuric acid.
Ternary Acids
Definition: Acids whose anion is an oxoanion, named based on whether the anion ends in -ate or -ite.
Important details:
-ATE ending: replace -ate with -ic + "acid." Formula: root word + ic acid.
-ITE ending: replace -ite with -ous + "acid." Formula: root word + ous acid.
From the material: H2CO3 = carbonic acid; HNO2 = nitrous acid; H3PO3 = phosphorous acid; H2SO4 = sulfuric acid.
Example: HClO3 (from chlorate, ClO3−) is named chloric acid.
More Examples
Writing Ionic Formulas — Combining Ba²⁺ and Cl⁻ by crisscrossing charges gives BaCl2, since barium's +2 charge becomes the subscript for chlorine and chlorine's 1 (implied) charge becomes barium's subscript.
Molecular Compound Naming — SiO2 (silicon dioxide) uses the Greek prefix "di" to show 2 oxygen atoms are bonded to 1 silicon atom.
Naming Acids — H2SO3, which comes from the sulfite ion (SO3²⁻, -ite ending), is named sulfurous acid, following the -ite → -ous acid rule.
Naming Compounds with Multiple Cation Charges — CuO is named copper(II) oxide (or cupric oxide), since the oxide ion's -2 charge tells us the copper must be +2 to balance it.
Chemical Bonds and Properties of Compounds
Introduction
Chemical Bond
Definition: The attractive force that holds two or more atoms together to form simple or complex units, formed because of interactions between electrons in the combined atoms.
Important details:
Atoms rarely exist on their own in nature (the exceptions being the noble gases: helium, neon, argon).
This interaction can form ionic, covalent, or metallic bonds; this reviewer covers only ionic and covalent bonding.
Chemical bonding occurs because it results in a total energy that is lower than any alternative arrangement.
Example: The oxygen we breathe (O2) exists as two oxygen atoms held together by a chemical bond, rather than as separate, individual oxygen atoms.
Chemical Bonding
Fundamental Concepts in Bonding
Valence Electrons in Bonding
Definition: The key concept that only the electrons in the valence shell (valence electrons) are involved when atoms react to form chemical bonds.
Important details:
Only the electronic configuration in the valence shell is important in bond formation.
Certain arrangements of electrons are more stable than others — for example, the electron configuration of the noble gases.
Example: Only the outermost electron of a sodium atom (its single valence electron) participates in forming a bond with chlorine — its inner, filled shells stay untouched.
Lewis Symbol for Atoms
Lewis Symbol (Lewis Electron Dot Symbol, LEDS)
Definition: An abbreviated representation of the structure of an atom, showing only its valence electrons as dots surrounding the chemical symbol.
Important details:
The group A number of representative elements corresponds to their number of valence electrons and, consequently, the number of dots.
Structure (Lewis symbols for representative elements, from the material):
Group: 1A 2A 3A 4A 5A 6A 7A 8A
Valence e⁻: 1 2 3 4 5 6 7 8(except He)
Period 1: H· He:
Period 2: Li· Be: B· ·C· ·N: ·O: :F: :Ne:
Period 3: Na· Mg: Al· ·Si· ·P: ·S: :Cl: :Ar:
Example: Magnesium (group 2A) is represented as Mg with two dots, since it has 2 valence electrons.
The Octet Rule and Bond Formation
Octet Rule
Definition: The stability of an atom can be achieved by having 8 electrons in its outermost energy level.
Important details:
An octet can be considered as four pairs of valence electrons arranged around the atom (as seen in the group VIIIA elements).
To attain stable configuration, atoms enter chemical combinations where they give away, accept, or share valence electrons to achieve the same electron arrangement as the nearest noble gas.
Noble gases have full valence shells (8 electrons, except helium's 2), making them stable, inert, and unreactive; other atoms with incomplete valence shells are more likely to bond.
Example: A chlorine atom, which has 7 valence electrons, bonds with another atom to gain 1 more electron and complete its octet.
Primary Types of Bonding
Covalent Bonding
Definition: A type of bonding where an electron is shared, spending its time moving in the electron orbitals around both atoms.
Important details:
The compound formed is called a covalent or molecular compound.
Example: Two hydrogen atoms sharing their single electrons to form H2 gas is covalent bonding.
Ionic Bonding
Definition: A type of bonding where an electron is transferred from one atom to another — one atom gains an electron while the other loses one.
Important details:
This results in positively and negatively charged particles called ions.
The compound formed is called an ionic compound; an ionic bond is also called an electrovalent bond.
Example: A sodium atom transferring its one valence electron to a chlorine atom, forming Na+ and Cl− ions, is ionic bonding.
Characteristics of Ionic and Covalent Compounds
Structural Units of Ionic vs. Covalent Compounds
Definition: The basic building block of each compound type — a molecule for covalent compounds, and a formula unit for ionic compounds.
Important details:
A formula unit represents the lowest whole-number ratio of cation to anion; ionic compounds do not form molecules.
Lattice Energy
Definition: The amount of energy needed to completely separate the positive and negative ions of an ionic compound apart.
Important details:
The more energy required, the stronger the attraction between the ions, and the stronger the ionic bond.
Two factors affect ionic bond strength: the magnitude of the charges (greater charge = stronger bond) and the size of the ion (smaller ion = stronger bond, since it allows ions to get closer together).
For cations, ionic size decreases as charge increases; for anions, ionic size increases as the magnitude of negative charge increases.
Example: MgO has a higher lattice energy than NaCl because magnesium and oxide ions have greater charges (+2 and −2) than sodium and chloride ions (+1 and −1).
General Properties of Ionic Compounds
Definition: The characteristic physical properties shared by ionic compounds.
Important details:
High melting and boiling points (typically 400°C+) — due to strong electrostatic attraction between ions; larger lattice energy = higher melting point.
Solid (crystalline) at room temperature — the electrostatic attraction draws ions into a solid lattice arrangement.
Hard and brittle solids — striking a crystal shatters it since a lot of energy is needed to shift ions apart.
Conduct electricity in solution, molten, or liquid state — solid ions are locked in place (poor conductors), but ions become mobile when melted or dissolved.
Most are soluble in water (though not all ionic solids are soluble).
Example: Table salt (NaCl) doesn't conduct electricity as a solid crystal, but a saltwater solution does conduct electricity because the Na+ and Cl− ions become mobile.
General Properties of Covalent Compounds
Definition: The characteristic physical properties shared by covalent (molecular) compounds.
Important details:
Generally lower melting and boiling points (typically 300°C or less).
Solid, liquid, or gas at room temperature.
Brittle and weak, or soft and waxy, solids.
Mostly do not conduct electricity, since they remain as molecules even when dissolved in water (though some molecular ionizing compounds do conduct, either strongly or weakly).
Solubility depends on molecular structure; often soluble in nonpolar solvents like hexane and carbon tetrachloride, and less soluble in polar solvents like water.
Example: Sugar (a covalent compound) dissolves in water but the resulting solution does not conduct electricity, unlike a salt solution.
The Ionic Bond
Nature of Ionic Bonding
Ionic Bond
Definition: A bond formed through electrostatic attraction between two oppositely charged ions, created when one or more electrons transfer from the valence shell of one atom to another.
Important details:
If an atom gains electron(s), it becomes a negatively charged anion. If it loses electron(s), it becomes a positively charged cation.
Both resulting ions become isoelectronic to a noble gas (matching its electron configuration).
Four key points: matter is conserved (no atom loses electrons without another gaining the same number); metals lose electrons to nonmetals; metals are cation formers while nonmetals are anion formers; atoms become isoelectronic to a noble gas when they gain or lose electrons.
Structure (from the material — sodium and chlorine ion formation):
Atom: 11Na 17Cl
Electron config: 1s²2s²2p⁶3s¹ 1s²2s²2p⁶3s²3p⁵
Lewis symbol: Na· :Cl:
¨¨
Ion: 11Na⁺ 17Cl⁻
Electron config: 1s²2s²2p⁶ (Ne) 1s²2s²2p⁶3s²3p⁶ (Ar)
Lewis symbol: [Na]⁺ [:Cl:]⁻
¨¨
Example: A calcium atom (2 valence electrons) transferring both electrons to two fluorine atoms (each needing 1) forms CaF2, with Ca²⁺ and two F⁻ ions.
Ionic Energetics: Ionization Energy and Electron Affinity
Ionization Energy (IE)
Definition: The amount of energy required to remove the most loosely held electron from an isolated gaseous atom or ion in its ground state; the energy change when an atom loses an electron to form a cation.
Formula:
Important details:
IE increases from left to right across a period (atoms get smaller, valence electrons more strongly attracted to the nucleus) and decreases from top to bottom in a group (atoms get larger, valence electrons less strongly attracted).
Elements with low IEs easily lose electrons to form cations — this is why metals are cation formers.
Example: Sodium has a lower ionization energy than chlorine, making it much easier for sodium to lose an electron and form a cation.
Electron Affinity (EA)
Definition: The amount of energy released when an electron is added to an isolated atom or neutral atom in its ground state; the energy change when an atom gains an electron to form an anion.
Formula:
Important details:
EA becomes increasingly negative from left to right across a period (atoms become smaller, easier for incoming electrons to be attracted) and less negative going down a group (atoms become larger, harder for incoming electrons to be added).
Non-metals with lower EA have a greater tendency to accept electrons and form an anion.
An ionic bond is formed when electron(s) from a metal (low IE) transfer to a nonmetal (low EA).
Example: Fluorine has a very negative electron affinity, meaning a large amount of energy is released when it gains an electron, making it a strong anion former.
Lewis Structure for Ionic Compounds
Lewis Structure (for Ionic Compounds)
Definition: A combination of Lewis symbols that represents the transfer of electrons in ionic bond formation, showing the number of valence electrons and how they're involved in bonding.
Important details:
Ionic compound formation requires a metal (donates electrons) and a nonmetal (accepts electrons); the direction of the arrow in the diagram must point from the metal to the nonmetal.
Structure (from the material — formation of NaCl):
Na ⟶(+) ·Cl: ⟶ [Na]⁺ [:Cl:]⁻ ⟶ NaCl
¨¨ ¨¨
Structure (from the material — formation of Na2O, where 2 sodium atoms each give 1 electron to oxygen):
Na
\
·O: ⟶ [Na]⁺
/ [Na]⁺ [:Ö:]²⁻ ⟶ Na2O
Na (or written as: 2[Na]⁺, [:Ö:]²⁻)
Example: For MgCl2, the Lewis structure would show one Mg atom transferring one electron each to two separate Cl atoms, giving [Mg]²⁺ and two [Cl]⁻ ions.
The Covalent Bond
Electronegativity and the Nature of Covalent Bonding
Electronegativity (EN)
Definition: The ability of an atom to attract electrons toward itself when chemically combined with another atom.
Important details:
For representative elements, EN usually increases from left to right across a period and decreases from top to bottom within a group (except noble gases).
Fluorine is the most electronegative element. Electronegativity values are unitless.
The most common scale for quantifying electronegativity is the Pauling scale (other scales: Mulliken, Allred-Rochow).
Important details:
Covalent bonding occurs between two atoms with the same or nearly the same electronegativity — nonmetals and metalloids have high electron-attracting power, so instead of one atom giving up electrons, they share valence electrons to reach 8 electrons in the outermost shell.
Covalent bonds form as valence orbitals overlap when atoms come closer together, allowing electron sharing that attracts both nuclei and forms a bond.
Example: Two chlorine atoms, both with the same electronegativity, share electrons equally to form Cl2 rather than one atom taking electrons from the other.
Lewis Structures for Molecular Compounds
Bond Pair and Lone Pair
Definition: The two types of electron pairs shown in a Lewis structure.
Important details:
Bond pairs are pairs of valence electrons shared between atoms in a covalent bond, represented with a line or dash.
Lone pairs are pairs of valence electrons on an atom not involved in bonding (also called unshared or nonbonding electron pairs).
Structure (from the material):
H—H H—F: :F—F: :Br—F:
¨¨ ¨¨ ¨¨ ¨¨
(lone pair on F) (bond pair between Br and F)
Single Covalent Bond
Definition: A covalent bond constituting one shared pair of electrons.
Important details:
Examples: H2, Br2, Cl2, HI.
Structure (from the material — formation of HCl):
H· ·Cl: → H:Cl: or H—Cl:
¨¨ ¨¨ ¨¨
Example: HBr is a single covalent bond, with hydrogen and bromine each contributing one electron to the shared pair.
Double Covalent Bond
Definition: A covalent bond constituting two shared pairs of electrons.
Important details:
Examples: O2, CO2.
Structure (from the material — formation of O2):
:O::O: or :O=O:
¨¨ ¨¨ ¨¨ ¨¨
Example: In CO2, each C=O bond is a double covalent bond, made of two shared electron pairs.
Triple Covalent Bond
Definition: A covalent bond constituting three shared pairs of electrons.
Important details:
Example: N2.
Structure (from the material — formation of N2):
:N⋮⋮N: or :N≡N:
Example: The strong triple bond in N2 (N≡N) is one reason nitrogen gas is so chemically stable and unreactive.
Exceptions to the Octet Rule
Definition: Categories of molecules where the octet rule does not apply.
Important details:
Incomplete Octet — involves smaller central atoms that hold fewer than 8 electrons, such as boron (maximum of 6 electrons).
Expanded Octet — involves larger central atoms (third period and below) that can hold more than 8 electrons because they have empty d orbitals; commonly seen with S and P.
Molecules with an odd number of electrons — such as NO and NO2, sometimes called radicals; species that don't obey the octet rule are relatively reactive.
Structure (from the material — BH3, incomplete octet):
H
|
B
/ \
H H
Structure (from the material — SF4, expanded octet):
:F:F:
\ | /
:F — S — F:
Example: PCl5 is an expanded-octet molecule, since phosphorus (a third-period element) can hold more than 8 electrons around it.
Steps in Writing Lewis Structures of Polyatomic Molecules
Definition: The systematic process for constructing a Lewis structure.
Important details (in order):
Determine the central atom — the bigger atom, with higher covalency (can form more bonds), and less electronegative.
Atom placement — write the central atom at the center, placing other atoms around it (H and F are usually attached, outer atoms).
Count total valence electrons (TVe) — add all A-group numbers of the atoms in the formula.
Determine bonding and nonbonding electrons (trial-and-error): a. Assume all bonds are single, drawing one line between the central atom and attached atoms. b. Count remaining electrons: TRe = TVe − bonding electrons. c. Distribute remaining electrons in pairs, starting on the attached atoms (giving them 8 electrons, except H), before the central atom. d. Check if the central atom obeys the octet rule; if not, form multiple bonds and repeat.
Step-by-step example from the material (PBr3): Central atom = P. TVe = 1(5) + 3(7) = 26. Assuming single bonds uses 6 electrons, so TRe = 26 − 6 = 20. Distributing gives each Br 3 lone pairs; central atom P obeys the octet rule.
Structure (from the material — Lewis structure of PBr3):
Br
|
Br — P — Br
(lone pair on P not shown here but present)
Step-by-step example from the material (SO2): Central atom = S. TVe = 2(6) + 1(6) = 18. Single bonds use 4 electrons, TRe = 18 − 4 = 14; central atom does not obey octet, so one lone pair forms a double bond, giving TRe = 18 − 6 = 12.
Structure (from the material — Lewis structure of SO2):
..
O — S = O or O = S — O
..
Step-by-step example from the material (SF4): Central atom = S. TVe = 1(6) + 4(7) = 34. Single bonds use 8 electrons, TRe = 34 − 8 = 26. Central atom does not obey the octet rule, but since S is a larger atom (period 3), it can accommodate more than 8 electrons (expanded octet).
Example: Following these same steps for CF4 gives a central carbon atom bonded singly to four fluorine atoms, each fluorine holding 3 lone pairs, with carbon obeying the octet rule (4 bond pairs = 8 electrons around C).
VSEPR Theory and Molecular Geometry
VSEPR Theory (Valence Shell Electron Pair Repulsion Theory)
Definition: The simplest bond model used to predict molecular shapes, based on the assumption that all valence electron pairs around a central atom (bonding or nonbonding) repel one another and stay as far apart as possible.
Important details:
The geometry of a molecule affects its physical and chemical properties and reactivity; it can be predicted even without experimentation.
Electronic geometry considers both bonding pairs and lone pairs on the central atom.
Molecular geometry only considers the arrangement of bonding (shared) electrons, ignoring lone pairs.
Important details — Key Ideas of VSEPR Theory:
Electron pairs stay far apart from each other to minimize repulsions.
Molecular shape is determined by the number of bond pairs and lone pairs around the central atom.
Multiple bonds (double or triple) are treated as if they were single bonds when predicting shape.
Lone pairs occupy more space than bond pairs (repelling effect: LP-LP > LP-BP > BP-BP).
Important details — Rules for Applying VSEPR Theory:
Draw the Lewis structure of the molecule.
Count the number of atoms bonded to the central atom, X.
Count the number of lone pairs on the central atom, E.
Add the total bonding pairs and lone pairs (X + E), then use the summary table to determine shape.
Summary Table of Geometries of Molecules
# atoms bonded (X) | # lone pairs (E) | Total (X+E) | VSEPR formula | Electronic Geometry | Molecular Geometry |
|---|---|---|---|---|---|
2 | 0 | 2 | AX2 | linear | linear |
3 | 0 | 3 | AX3 | trigonal planar | trigonal planar |
2 | 1 | 3 | AX2E | trigonal planar | V-shaped/bent |
4 | 0 | 4 | AX4 | tetrahedral | tetrahedral |
3 | 1 | 4 | AX3E | tetrahedral | trigonal pyramidal |
2 | 2 | 4 | AX2E2 | tetrahedral | V-shaped/bent |
5 | 0 | 5 | AX5 | trigonal bipyramidal | trigonal bipyramidal |
4 | 1 | 5 | AX4E | trigonal bipyramidal | see-saw |
3 | 2 | 5 | AX3E2 | trigonal bipyramidal | T-shaped |
2 | 3 | 5 | AX2E3 | trigonal bipyramidal | linear |
6 | 0 | 6 | AX6 | octahedral | octahedral |
5 | 1 | 6 | AX5E | octahedral | square pyramidal |
4 | 2 | 6 | AX4E2 | octahedral | square planar |
Structure (worked geometry examples from the material):
1.) CF4 — X=4, E=0, total 4 pairs → electronic: tetrahedral, molecular: tetrahedral
F
|
F — C — F
|
F
2.) PBr3 — X=3, E=1, total 4 pairs → electronic: tetrahedral, molecular: trigonal pyramidal
..
P
/ \
Br Br
\
Br
3.) SO2 — X=2, E=1, total 3 pairs → electronic: trigonal planar, molecular: bent
..
S
/ \
O O
4.) XeF4 — X=4, E=2, total 6 pairs → electronic: octahedral, molecular: square planar
F F
\ /
Xe
/ \
F F
(2 lone pairs positioned opposite each other)
Important details:
If there is one lone pair on the central atom, it can be placed anywhere in a linear, trigonal planar, tetrahedral, or octahedral molecule (equivalent positions); but for trigonal bipyramidal, the lone pair must be placed in the equatorial position (least crowded, 120° bond angle).
If there are two lone pairs on the central atom, for tetrahedral geometry they can be placed anywhere; for trigonal bipyramidal, they must be in 2 equatorial positions; for octahedral, they must be placed across from each other.
Example: NH3 (ammonia) has 3 bonded atoms and 1 lone pair on nitrogen (X=3, E=1), giving it a tetrahedral electronic geometry but a trigonal pyramidal molecular shape, just like PBr3.
Polarity of Bonds and Molecules
Bond Polarity
Definition: The concept describing how evenly or unevenly electrons are shared or distributed between two bonded atoms, directly related to the difference in electronegativities (ΔEN) between them.
Important details:
As ΔEN increases, bond polarity increases (the bond becomes more polar).
Take the absolute value of the difference; only the magnitude of ΔEN matters.
Nonpolar Covalent Bond
Definition: A bond formed when bonding electrons are shared equally between atoms; occurs between atoms with the same or a slight difference in electronegativities (ΔEN between 0.00–0.40).
From the material: F–F in F2 = 0.00; C–H in CH4 = 0.35.
Example: The Br–Br bond in Br2 is nonpolar since both atoms have identical electronegativity.
Polar Covalent Bond
Definition: A bond formed when there is unequal sharing of bonding electrons between two atoms; occurs when ΔEN is between 0.41–2.00.
From the material: O–H in H2O = 1.24; H–F = 1.78.
Example: The O–H bond in water is polar, since oxygen is significantly more electronegative than hydrogen.
Ionic Bond (in terms of ΔEN)
Definition: A bond classified as ionic when ΔEN between two bonded atoms is ≥ 1.70 and is shared between a metal and a nonmetal.
Important details:
No bond is truly 100% ionic, but a ΔEN of 2.0 or greater is usually classified as predominantly ionic.
Example: The Na–Cl bond in table salt, with a large ΔEN between a metal and nonmetal, is classified as ionic.
Dipole
Definition: A separation of charge created when bonding electrons are not shared equally, producing a partial positive charge (δ+) on the less electronegative atom and a partial negative charge (δ−) on the more electronegative atom.
Important details:
The charges are only partial because there is no complete electron transfer, only uneven distribution — a polar covalent bond lies between a purely covalent (nonpolar) bond and an ionic bond.
Structure (from the material — dipole in HCl):
δ+ δ-
H ⟶⟶⟶⟶ Cl
Example: In H2O, both O–H bonds are dipoles, with oxygen carrying the partial negative charge and each hydrogen carrying a partial positive charge.
Bond Dipole Moment
Definition: The measure that occurs when there is a separation of charge, represented as a vector (crossed arrow) pointing toward the more electronegative atom.
Important details:
Can occur in ionic bonds or covalent bonds. The larger the ΔEN, the longer the dipole moment.
Example: The dipole moment arrow in the C–H bond of methane points slightly toward carbon, since carbon is slightly more electronegative than hydrogen.
Molecular Dipole Moment / Molecular Polarity
Definition: The unequal distribution of charge across an entire molecule, determined by both the polarity of its bonds and its overall shape.
Important details:
If electrons are shared unevenly across the whole molecule, it is polar (has a partially positive end and a partially negative end).
If electrons are shared evenly, it is nonpolar.
For diatomic molecules, polarity depends only on the bond type. For polyatomic molecules, polarity depends on both individual bond polarity and molecular shape; the presence of lone pairs on the central atom also affects polarity.
Important details — A molecule is POLAR if:
One or more bonded atoms differ from each other, or at least one polar bond is present.
The molecule has one slightly positive end and one slightly negative end.
The molecule contains one lone pair (or lone pairs that are not symmetrically arranged so their bond dipoles don't cancel).
Important details — A molecule is NONPOLAR if:
All bonded atoms are the same.
All bonded atoms are symmetrically arranged around the central atom.
The bonded atoms have the same charges (or lone pairs, if present, are symmetrically arranged so their dipoles cancel).
Step-by-step examples from the material:
CF4 — all bonded atoms are the same and symmetrically arranged, bond dipoles cancel → nonpolar.
PBr3 — all bonded atoms are the same but not symmetrically arranged (1 lone pair on central atom), bond dipoles don't cancel → polar.
SO2 — all bonded atoms are the same but not symmetrically arranged (1 lone pair) → polar.
XeF4 — all bonded atoms are the same; the 2 lone pairs are symmetrically arranged (opposite each other), so dipoles cancel → nonpolar.
Example: CO2 is a linear, symmetric molecule with two identical polar C=O bonds pointing in opposite directions, so the bond dipoles cancel out, making CO2 nonpolar overall — even though each individual bond is polar.
More Examples
Octet Rule — A magnesium atom (2 valence electrons) loses both electrons to a nonmetal to achieve the same stable electron configuration as neon.
Ionic vs. Covalent Compound Properties — Table salt (NaCl) needs to be heated to over 800°C to melt, while sugar (a covalent compound) melts at a much lower temperature around 186°C.
Lewis Structures — Drawing the Lewis structure for water (H2O) shows oxygen with 2 bond pairs (to each hydrogen) and 2 lone pairs, giving it a bent molecular shape.
VSEPR Theory — Methane (CH4), with 4 bonded atoms and no lone pairs on carbon, has a tetrahedral molecular geometry, just like CF4.
Molecular Polarity — Even though the individual C=O bonds in CO2 are polar, the molecule's linear, symmetric shape makes the overall molecule nonpolar, similar to how CF4's bond dipoles cancel out.
Intermolecular Forces and Properties of Substances
Introduction
Intermolecular Forces of Attraction (IMFA)
Definition: The attractive forces between particles of a substance — described as invisible, "sticky" forces acting between neighboring particles.
Important details:
The strength of these forces determines how strongly particles hold together, thereby determining whether a substance is solid, liquid, or gas.
IMFA directly affects physical properties of substances, including surface tension, viscosity, vapor pressure, boiling point, and molar heat of vaporization.
From the material: Everyday observations explained by IMFA include why water beads up instead of spreading, why water flows faster than honey, why rubbing alcohol evaporates faster than water, and why water requires so much heat energy to boil.
Example: A drop of nail polish remover (acetone) disappearing within seconds while a grain of table salt does not is explained by the difference in the strength of their intermolecular/interparticle forces.
Intramolecular vs. Intermolecular Forces
Intramolecular Forces
Definition: The forces that hold atoms together within a molecule or ions together within an ionic compound — commonly called chemical bonds.
Example: The O–H bonds within a single water molecule are intramolecular forces.
Intermolecular (Interparticle) Forces
Definition: The attractive forces existing between particles (molecules, atoms, or ions) of the same substance or of different substances.
Important details:
Chemical bonds (intramolecular forces) are much stronger than intermolecular forces. When water boils, it is the intermolecular force that is overcome, not the intramolecular O–H bond.
It is the intermolecular force between molecules, not the intramolecular bond, that contributes to properties like water's high boiling point.
IMFA plays a significant role in determining whether substances exist as gas, liquid, or solid, and affects physical properties like melting point, boiling point, and viscosity.
Example: The attraction between neighboring H2O molecules in a glass of water is an intermolecular force, distinct from the intramolecular O–H bonds inside each molecule.
Types of Intermolecular Forces of Attraction (IMFA)
Common Feature of All IMFA Types
Definition: Regardless of type, all IMFA share a common feature: the attraction of opposite charges, whether full charges, partial charges, or induced charges.
Ion-Ion Interaction
Ion-Ion Interaction
Definition: The intermolecular force exhibited by ionic compounds, involving the attraction of the cation from one formula unit and the anion from another.
Important details:
Its energy of attraction is very high because it involves a full positive and full negative ion; most ionic compounds have high melting and boiling points.
In most cases, the energy of attraction between ions increases as their charge increases.
From the material: NaCl has a melting point of 801°C and a boiling point of 1,413°C.
Example: The strong ion-ion attraction between K+ and Br− ions in solid KBr gives it a high melting point.
Ion-Dipole Interaction
Ion-Dipole Interaction
Definition: The intermolecular force occurring between an ionic compound and a polar molecule, involving attraction between the full charge of the ion and the opposite partial charge of a polar molecule.
Important details:
The cations of the ionic compound are attracted to the partial negative end of the dipole; the anions are attracted to the partial positive end.
This interaction is significant in dissolving ionic compounds in polar liquids like water (e.g., aqueous NaCl), and in hydrated salts like cupric sulfate pentahydrate (CuSO4·5H2O).
Example: When table salt dissolves in water, the Na+ ions are surrounded by the partial negative (oxygen) end of water molecules, an example of ion-dipole interaction.
Dipole-Dipole Interaction
Dipole-Dipole Interaction
Definition: Also called permanent dipole-dipole interaction, this occurs in polar molecules due to attraction between the partially positive end of one polar molecule and the partially negative end of another.
Important details:
Weaker than ion-dipole interaction. Its strength is inversely proportional to the distance between molecules — closer molecules have stronger interaction.
The strength also increases with increasing polarity: for two polar molecules with similar molar mass and size, the more polar molecule has a stronger dipole-dipole force.
Example: The attraction between the partial positive hydrogen of one HCl molecule and the partial negative chlorine of a neighboring HCl molecule is a dipole-dipole interaction.
Hydrogen Bonding
Hydrogen Bonding (H-Bonding)
Definition: A special, especially strong type of dipole-dipole interaction that occurs among polar covalent molecules containing hydrogen bonded to one of three small, highly electronegative elements: F, O, or N.
Important details:
H-bonding is 4 to 5 times stronger than an ordinary dipole-dipole interaction.
It results from the attraction between the partial positive H-atom and a lone pair of electrons on the F, O, or N atom of another molecule.
It is stronger than ordinary dipole-dipole interactions because: (1) F, O, and N are highly electronegative, making the H bond highly polar; (2) H is so small that the electrons of F, O, or N can be strongly attracted to its nucleus.
Molecules that exhibit H-bonding have unusually high melting and boiling points compared to compounds of similar molar mass and geometry — water has the highest boiling point among comparable substances largely due to H-bonding.
H-bonding also plays an important role in biological systems, such as helping stabilize the structure of DNA.
Example: The strong hydrogen bonds between ammonia (NH3) molecules, since nitrogen is bonded to hydrogen, give ammonia a higher boiling point than would otherwise be expected for its small size.
London Dispersion
London Dispersion Force
Definition: An intermolecular force where the motion of electrons in an atom or molecule creates an instantaneous (momentary) dipole moment, discovered by physicist Fritz London.
Important details:
Since electrons are always moving, at a given moment they may be more concentrated on one side, creating a temporary dipole. This temporary dipole can induce a similar dipole in a neighboring molecule, causing attraction.
London dispersion is present in all types of molecules in condensed phases (solid and liquid) — without it, substances like CO2 or I2 couldn't condense or solidify.
It is the only type of IMFA present in nonpolar molecules and nonpolar atoms (like noble gases).
It is weak in small molecules but generally stronger in larger, more polarizable molecules with more electrons; it is also short-ranged (effective over short distances) and increases in strength with increasing molecular weight and size.
Important details — In summary:
No dipole-dipole force exists between nonpolar atoms/molecules, since they lack a permanent dipole moment.
In slightly polar or nonpolar covalent compounds or monatomic nonmetals, dispersion forces (though weak) are still the strongest forces present.
For large molecules, London dispersion becomes significant, since large molecules/atoms are highly polarizable (contain many electrons far from the nucleus).
Example: Two neighboring nonpolar CO2 molecules attracting each other through momentarily-induced dipoles is an example of London dispersion, allowing CO2 to become dry ice at low enough temperatures.
Summary Table of the Five Intermolecular (Interparticle) Forces
Type of IMFA | Type of Interaction | Occurrence |
|---|---|---|
Ion-ion | An ion is attracted to a nearby oppositely charged ion | Ionic compounds |
Ion-dipole | Attraction of ions (from an ionic compound) and the opposite partial charge of a polar molecule | Polar molecule and ionic compound |
H-Bonding | One polar molecule containing H bonded to an electronegative element is attracted to the lone pair of another polar molecule | Polar molecules with a lone pair of electrons and H bonded to F, O, N |
Dipole-dipole | Opposite dipole ends of polar molecules attract one another | Polar molecules |
London Dispersion | A temporary dipole in one molecule induces a temporary dipole in a nearby molecule and is attracted to it | All polar molecules, nonpolar molecules, and noble gases |
Determining the Dominant IMFA (Sample Exercises from the material)
Compound | Interacting Particles | Dominant IMFA |
|---|---|---|
SO2 | Polar molecule | Dipole-dipole |
NaCl | Ionic compound/ions | Ion-ion |
CO2 | Nonpolar molecule | London dispersion |
H2O | Polar molecule | H-bonding |
NaCl in H2O | NaCl (ionic) + H2O (polar) | Ion-dipole |
Example: For the compound HF, since hydrogen is bonded to the highly electronegative fluorine, its dominant IMFA is hydrogen bonding — the same logic used to determine the dominant force for H2O.
IMFA and the Physical Properties of Substances
Surface Tension
Surface Tension
Definition: A property of a liquid that allows its surface to resist an external force, caused by strong cohesion (cohesive force) at the surface of liquids.
Important details:
Cohesive force is an intermolecular attractive force that binds similar molecules to one another.
Molecules at the liquid's surface experience an imbalance of attractive forces (fewer molecules above them to interact with), so they are pulled inward, reducing the surface. Inner molecules don't experience this because forces cancel out in all directions.
Surface tension is directly proportional to the strength of intermolecular force, and decreases with increasing temperature.
Water shows distinctly high surface tension due to strong hydrogen bonding, allowing a paper clip to float or a water strider to walk on its surface.
Adhesion (adhesive force)
Definition: The intermolecular attractive force between different types of molecules, such as between water and a waxed surface, or between a liquid and its container.
From the material: Water "beads up" on a waxed surface because water's cohesive force is greater than the adhesive force between water and the wax.
Meniscus and Capillary Action
Definition: The curved surface a liquid forms in a narrow container, shaped by the balance between adhesive and cohesive forces.
Important details:
Water's meniscus curves upward (concave/U-shaped) because its adhesive force (attraction to the glass) is greater than its cohesive force, causing it to "wet" the glass — this also causes capillary action (spontaneous rise of liquid in narrow tubes).
Mercury's meniscus curves downward (dome-shaped/convex) because its cohesive force is stronger than its adhesive force to glass, so it does not wet the glass and does not show capillarity.
Water in nonpolar plastic containers forms a flat meniscus, since neither adhesive nor cohesive forces are strong enough to dominate.
Example: A dry paper towel placed with one end touching a puddle of water will draw water upward through its fibers via capillary action, driven by water's strong adhesive force.
Viscosity
Viscosity
Definition: A measure of the resistance of a liquid to flow, related to the ease with which individual molecules of the liquid can move with respect to one another; the opposite of fluidity.
Important details:
The stronger the intermolecular forces, the greater the viscosity and the more slowly the liquid flows.
Viscosity is inversely proportional to temperature — as temperature increases, viscosity decreases, since the increase in kinetic energy overcomes the forces holding molecules together.
From the material (Sample Exercise): Glycerol is more viscous than ethanol, since although both molecules exhibit hydrogen bonding, glycerol has more sites for H-bonding.
From the material (Practical Application): Engine oil viscosity is graded using a code like "SAE 5W-30," where the number before "W" (winter) describes viscosity at low temperatures (lower = thickens less in cold weather) and the number after describes hot weather viscosity/resistance to thinning (higher = thicker at high temperatures).
Example: Honey is more viscous than water because honey's molecules have stronger intermolecular attractions, making it flow much more slowly.
Vapor Pressure and Boiling Point
Vapor Pressure
Definition: The pressure exerted by a vapor in equilibrium with a solid or a liquid at a given temperature.
Important details:
Equilibrium occurs when evaporation happens as fast as condensation, so there's no net change in the amount of vapor in a closed container.
Vapor pressure is also a measure of a liquid or solid's escaping tendency; substances with high vapor pressure readily evaporate and are said to be volatile.
The greater the attractive force between particles, the lower the escaping tendency and vapor pressure. Increasing temperature increases kinetic energy, which overcomes attractive forces and increases vapor pressure.
Boiling Point
Definition: The temperature at which a liquid's vapor pressure equals the external (atmospheric) pressure acting on its surface, causing the liquid to boil.
Important details:
The normal boiling point is the temperature at which a liquid boils under a pressure of 760 mm Hg (one atmosphere).
Liquids with weak attractive forces have high vapor pressure and boil at lower temperatures; those with strong attractive forces boil at higher temperatures.
A liquid's temperature stays constant at its boiling point, because added heat energy is used to overcome cohesive forces rather than raise the temperature further.
Step-by-step example from the material: Predicting increasing boiling points for H2S, H2O, CH4, H2, and KBr. Answer, from lowest to highest: H2 < CH4 < H2S < H2O < KBr. H2 and CH4 are both nonpolar, but CH4 is larger and more polarizable, so it has stronger dispersion forces and boils higher than H2. H2S and H2O are both polar, but H2O has stronger H-bonding than H2S's dipole-dipole interaction, so H2O boils higher. KBr, being ionic, boils at the highest temperature of all.
Example: Rubbing alcohol has a lower boiling point than water because its intermolecular forces, while including hydrogen bonding, are weaker overall, giving it higher vapor pressure and faster evaporation.
Molar Heat of Vaporization
Molar Heat of Vaporization
Definition: The amount of heat required to vaporize a specific amount of liquid at its boiling point with no change in temperature — the amount of energy needed to change a liquid into a vapor.
Important details:
Depends on the strength of the attractive forces in the liquid; liquids with stronger intermolecular forces have a higher heat of vaporization.
From the material: This is why a steam burn is more damaging to skin/tissue than a burn from boiling water at the same temperature — when steam condenses into liquid on the skin, a large amount of energy is released and absorbed by the skin.
Example: Water has a high molar heat of vaporization due to its strong hydrogen bonding, which is why sweating (evaporating water from skin) is such an effective way for the human body to cool down.
More Examples
Ion-Dipole Interaction — Adding potassium chloride (KCl) to water and watching it fully dissolve, with the K⁺ and Cl⁻ ions surrounded by water molecules, demonstrates ion-dipole interaction.
Hydrogen Bonding — Two ethanol molecules attracting each other through their O–H groups, similar to how water molecules attract each other, is an example of hydrogen bonding.
London Dispersion — Liquid nitrogen (N2) staying condensed at very low temperatures, despite N2 being a nonpolar molecule, is possible only because of London dispersion forces.
Surface Tension — A small insect resting on the surface of a pond without sinking is supported by the water's surface tension, caused by strong hydrogen bonding between water molecules.
Viscosity — Cold honey pouring much more slowly than warm honey demonstrates how lower temperature increases viscosity.
Vapor Pressure and Boiling Point — Water boils at a lower temperature at high mountain altitudes because the lower atmospheric pressure allows the liquid's vapor pressure to match it sooner.