Chemistry Notes

The Nature of Chemistry

  • Chemistry is the science of the composition, structure, properties, and reactions of matter.

  • It is also the science dealing with the composition of matter and the changes in composition that matter undergoes.

  • Chemistry is fundamental to many fields, including engineering, teaching, healthcare, and environmental science.

  • Studying chemistry helps develop problem-solving and communication skills.

  • Chemists observe nature, conduct experiments, and synthesize substances to advance knowledge and benefit humanity.

Thinking Like a Chemist

  • Chemists analyze the behavior of basic components by "looking inside" everyday objects.

  • The macroscopic picture is the overall view of an object or system.

  • The microscopic picture involves observing details at the atomic and molecular level.

  • Chemistry connects the microscopic world of molecules and the macroscopic world of everyday objects.

Enhanced Example 1.1

  • Given 8 oxygen atoms and 15 hydrogen atoms, you can make only 7 water molecules (H_2O) because each water molecule requires 2 hydrogen atoms, leaving one hydrogen atom leftover.

A Scientific Approach to Problem Solving

  • A logical approach is useful for solving daily problems

    • Define the problem by stating it clearly, including all known information. In science, this is called making an observation.

    • Propose possible solutions to the problem. In science, this is called making a hypothesis.

    • Decide which is the best way to proceed or solve the problem. In science, an experiment is performed.

  • A scientific approach to problem solving is worthwhile in all parts of life.

Chemistry in Action: Egyptians, the First Medicinal Chemists

  • Ancient Egyptians used lead in eyeliner for perceived health benefits, despite its toxicity.

  • Research suggests lead salts in the eyeliner may have triggered an immune response.

The Scientific Method

  • Scientists use the scientific method to solve problems:

    • Collect relevant facts or data through planned experimentation.

    • Analyze data to find trends.

    • Formulate a testable hypothesis.

    • Conduct additional experiments to test the hypothesis.

    • Modify the hypothesis as necessary.

  • A hypothesis is a tentative explanation of certain facts providing a basis for further experimentation.

  • A well-established hypothesis is often called a theory or model, explaining general principles with considerable supporting evidence.

  • Scientific laws are simple statements of natural phenomena with no known exceptions.

  • Science is dynamic; theories and models evolve with new experimental evidence.

  • Figure 1.2 illustrates the flow chart for the scientific method.

The Particulate Nature of Matter

  • The universe consists of matter and energy.

  • Matter is anything that has mass and occupies space.

  • Matter may be invisible, such as air.

  • Matter appears continuous to the macroscopic eye but is composed of discrete, tiny particles called atoms at the microscopic level.

Physical States of Matter

  • Matter exists in three physical states: solid, liquid, and gas.

  • A solid has a definite shape and volume, and its particles cling rigidly to one another.

  • Crystalline solids have particles arranged in regular, repeating, three-dimensional geometric patterns.

  • Amorphous solids (e.g., plastics, glass) lack a regular internal geometric pattern.

  • A liquid has a definite volume but not a definite shape, and its particles stick firmly but not rigidly.

  • A gas has indefinite volume and no fixed shape, with particles that move independently of one another.

  • Table 1.1 lists common materials in solid, liquid, and gaseous states.

  • Table 1.2 compares the properties of solids, liquids, and gases.

Classifying Matter

  • Matter can be classified as either homogeneous or heterogeneous.

  • A substance is a particular kind of matter with a definite, fixed composition (either an element or a compound).

  • Homogeneous matter is uniform in appearance and has the same properties throughout.

  • Heterogeneous matter consists of two or more physically distinct phases.

  • A phase is a homogeneous part of a system separated from other parts by physical boundaries.

  • A system is the body of matter under consideration.

  • A pure substance may exist as different phases in a heterogeneous system (e.g., ice floating in water).

  • A mixture is a material containing two or more substances and can be either heterogeneous or homogeneous.

  • Solutions are homogeneous mixtures.

  • Figure 1.7 illustrates the relationship of substances and mixtures.

Distinguishing Mixtures from Pure Substances

  • Mixtures always contain two or more substances in variable amounts, while pure substances have a definite composition by mass.

  • The components of a mixture do not lose their identities and can be separated by physical means, while elements in a compound lose their identities and can be separated only by chemical means.
    Example on creating homogeneous mixtures with salt and water, and heterogeneous mixtures with sulfur crystals and iron fillings.

Chapter 1 Review

  • Chemistry is the science dealing with matter and its changes.

  • The scientific method involves collecting facts, formulating hypotheses, planning experiments, and modifying hypotheses as necessary.

  • Matter has mass and occupies space and can exist as a solid, liquid, or gas.

  • Matter can be classified as a pure substance (element or compound) or a mixture (homogeneous or heterogeneous).

Measurements

Doing an experiment in chemistry is very much like cooking a meal in the kitchen. Measurements are an important part of the scientific process.

  • Quantitative observations require both a number and a unit.

Scientific Notation

  • Scientists use scientific notation to express very large or very small numbers.

  • A number in scientific notation is written as the product of a number between 1 and 10 multiplied by 10 raised to some power.

  • To write a number in scientific notation:

    • Move the decimal point in the original number so that it is located after the first nonzero digit.

    • Multiply this new number by 10 raised to the proper exponent (power). The proper exponent is equal to the number of places that the decimal point was moved.

    • The sign on the exponent indicates the direction the decimal was moved.

      • moved right {\rightarrow} negative exponent

      • moved left {\rightarrow} positive exponent
        Enhanced Example 2.1 - 2.3 shows the correct methods for implementing scientific notations.

Measurement and Uncertainty

  • Measurements always involve some degree of uncertainty.

  • It is important when recording a measurement to include all the digits that are known plus one digit that is estimated. This last estimated digit introduces some uncertainty.

Significant Figures

  • The digits used to express a measured quantity are known as significant figures.
    Enhanced Example 2.4 provides an example of estimated digits.

Rules for Counting Significant Figures

  • Nonzero digits: All nonzero digits are significant.

  • Exact numbers: Some numbers are exact and have an infinite number of significant figures.

    • Exact numbers occur in simple counting operations; when you count 25 dollars, you have exactly 25 dollars.

    • Defined numbers, such as 12 inches in 1 foot, 60 minutes in 1 hour, and 100 centimeters in 1 meter, are also considered to be exact numbers. Exact numbers have no uncertainty.

  • Zeros: A zero is significant:

    • Between nonzero digits: 205 has three significant figures (2, 0, 5)

    • At the end of a number that includes a decimal point:

  1. 500 has three significant figures (5, 0, 0)

  • A zero is not significant:

    • Before the first nonzero digit: These zeros are used to locate a decimal point: 0.0025 has two significant figures (2, 5)

    • At the end of a number without a decimal point: 1000 has one significant figure (1)
      *These rules should be memorized.
      *Enhanced Example 2.5 presents and explains a variety of examples that follows the rule for counting significant figures.*

Rounding Off Numbers

  • When digits are dropped from a number, the value of the last digit retained is determined by a process known as rounding off numbers.

Rules for Rounding Off

  • When the first digit after those you want to retain is 4 or less, that digit and all others to its right are dropped. The last digit retained is not changed.

  • When the first digit after those you want to retain is 5 or greater, that digit and all others to the right are dropped and the last digit retained is increased by one.

Significant Figures in Calculations

  • The results of a calculation based on measurements cannot be more precise than the least precise measurement.

Multiplication or Division

  • In calculations involving multiplication or division, the answer must contain the same number of significant figures as in the measurement that has the least number of significant figures.

Enhanced Example 2.6-2.7 presents a real world application of significant figures in multiplication and division.

Addition or Subtraction

  • The results of an addition or a subtraction must be expressed to the same precision as the least precise measurement. This means the result must be rounded to the same number of decimal places as the value with the fewest decimal places
    Enhanced Examples 2.8 - 2.10 uses a variety of problems implementing addition and subtraction with significant figures.

The Metric System

  • The metric system, or International System (SI, from Système International), is a decimal system of units for measurements of mass, length, time, and other physical quantities.

Common Prefixes and Numerical Values for SI Units

Table 2.1 shows the names, symbols, and numerical values of the common prefixes.

International System's Standard Units of Measurement

Table 2.2 shows how the abbreviations are determined.

Measurement of Length

  • The standard unit of length in the metric system is the meter (m).

  • One meter equals 10 decimeters, 100 centimeters, or 1000 millimeters.

  • A kilometer contains 1000 meters.

Metric Units of Length

Table 2.2 shows how many meters each unit is equivalent to.

Unit Conversions

  • One of the benefits of using the metric system is the ease with which we can convert from one unit to another.

  • To convert from one unit to another we must use a conversion factor. A conversion factor is a ratio of equivalent quantities.
    Table 2.4 displays length equivalent units.

Measurement of Mass

  • In science, the mass of an object is defined as the amount of matter in the object. Mass is measured on an instrument called a balance.

  • The weight of an object is a measure of the effect of gravity on the object. Weight is determined by using an instrument called a scale, which measures force against a spring.

  • The standard unit of mass in the SI system is the kilogram (kg) (equal to 1000 g).

  • A pound is equal to 453.6 g (0.4536 kg).
    Table 2.5 displays metric units of mass.

Measurement of Volume

  • Volume, as used here, is the amount of space occupied by matter.

  • The SI unit of volume is the cubic meter (m^3).

  • The liter (L) and the milliliter (mL) are the standard units of volume used in most chemical laboratories.

Dimensional Analysis: A Problem-Solving Method

  • Dimensional analysis converts one unit to another unit by using conversion factors.

Percent

  • Percent is defined as parts per hundred.

    ewline \text{Percent} = \frac{\text{Part}}{\text{Total}} \times 100

  • In chemistry, mass percent where
    \text{Mass Percent} = \frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100

Measurement of Temperature

  • Thermal energy is a form of energy associated with the motion of small particles of matter.

  • Temperature is a measure of the intensity of thermal energy, or how hot a system is, regardless of its size.

  • The term heat refers to the flow of energy due to a temperature difference.

  • The SI unit of temperature is the kelvin.

  • The common laboratory instrument for measuring temperature is a thermometer.

  • The temperature of a system can be expressed by several different scales.

    • Three commonly used temperature scales are Celsius, Kelvin (absolute), and Fahrenheit.
      \text{degrees Celsius} = ^\circ C
      \text{Kelvin (absolute)} = K
      \text{degrees Fahrenheit} = ^\circ F
      K = ^\circ C + 273.15
      ^\circ F = (1.8 \times ^\circ C) + 32

Density

  • Density (d) is the ratio of the mass of a substance to the volume occupied by that mass; it is the mass per unit of volume and is given by the equation
    Density = \frac{mass}{volume}

Chapter 2 Review

  • Quantitative observations consist of a number and a unit and are called measurements. Very large and very small numbers can be represented compactly by using scientific notation.
    *All measurements reflect some amount of uncertainty, which is indicated by the number of significant figures in the measurement. The significant figures include all those known with certainty plus one estimated digit.

  • The metric system uses factors of 10 and a set of standard units for measurements. Length in the metric system is measured by the standard unit of the meter. The standard unit for mass in the metric system is the kilogram. In chemistry, we often use the gram instead as we tend to work in smaller quantities. Volume is the amount of space occupied by matter.*The formula for percent is: \text{Percent = (part/total) * 100}* There are three commonly used temperature scales: Fahrenheit, Celsius, and Kelvin.
    *All formula of density is: Density = \frac{m}{V} where m=mass and V = volume

Elements
  • Element are a fundamental or elementary substance that cannot be broken down by chemical means to simpler substances.*
    Table 3.1 lists the abundant elements in the earth's crust, oceans, and the atmosphere.
    *Fourteen chemical symbols are single letter: C,H,N,O,F,I,P,S,U,Y,V,W. Twelve elements has a single letter as their symbol.Boron B; Carbon,C; Fluorine,F; Hydrogen,H; Iodine,I; Nitrogen,N; Oxygen,O; Phosphorus,P; Potassium,K; Sulfur,S; Uranium,U; Vanadium,V; Tungsten,W; Yttrium,Y.

Introduction to the periodic table.
  • Shows all the chemical elements and contains a great deal of useful information about them.*

  • Vertical column are labeled as families/groups .
    Transition Metals Metal, Nonmetals and Metalloid . Metals are has high luster, are good conductors of heat and electricity, are malleable (can be rolled or hammered into sheets), and are ductile (can be drawn into wires). Most metals have a high melting point and a high density. *TABLE 3.6 | Elements That Exist as Diatomic Molecules Element Symbol Molecular formula Normal state Hydrogen H H2 Colorless gas Nitrogen N N2 Colorless gas Oxygen O O2 Colorless gas Fluorine F F2 Pale Yellow gas Chlorine Cl Cl2 Greenish-yellow gas Bromine Br Br2 Reddish-brown liquid Iodine I I2

Compound and Formulas
  • A compound is distinct substance that contains two or more elements chemically combined in a definite proportion by mass.*

  • Molecule: is the smallest uncharged individual unit of a compound formed by the union of two or more atoms.*

  • An Ion is a positively or negatively charged atom or group of a atoms. If its positive it is a Cation Anion is negative charged. Chemical Formula shows the symbols and the ratio of the atoms of the elements in a compoundThe formula of a compound tells us which elements it is composed of and how many atoms of each element are present in a formula unit.
    A natural law is a summary of observed behavior. A model (theory) is an attempt to explain the observed behavior. There is the Dalton Atomic Model

Properties Of Substances
Physical And Chemical Changes
  • Different Types Of Energy

Chemistry In Action
  • Popping Popcorn Energy In The Real World

Types of Energy.
  • Heat

  • Light

  • Electrical

  • Sound

  • Motion

  • Energy Changes in the ecosystem.
    Chemical and Physical changes

Energy transformations and conservation

Energy Transformations:

  • Heat: Quantitative Measurement

    • Heat is a form of energy transfer that occurs between objects or systems due to a temperature difference.

    • It is a quantitative measurement of thermal energy, typically measured in units like joules (J) or calories (cal).

    • Heat can be transferred through three primary mechanisms:

      • Conduction: The transfer of heat through direct contact between particles or objects. It occurs when there is a temperature gradient within a material or between materials in contact.

      • Convection: The transfer of heat through the movement of fluids (liquids or gases). It occurs when warmer, less dense fluid rises, displacing cooler, denser fluid, creating a circulating current.

      • Radiation: The transfer of heat through electromagnetic waves. It does not require a medium and can occur in a vacuum. All objects emit thermal radiation, with the amount and wavelength depending on their temperature.

    Units of Heat
    • Joule (J): The SI unit of energy and heat. One joule is defined as the amount of energy required to exert a force of one newton over a distance of one meter.

      1 J = 1 N \cdot m

    • Calorie (cal): The amount of heat required to raise the temperature of 1 gram of water by 1 degree Celsius.

      1 cal = 4.184 J

    • Kilocalorie (kcal) or Calorie (Cal): Often used in nutrition, 1 kcal is the amount of heat required to raise the temperature of 1 kilogram of water by 1 degree Celsius.

      1 kcal = 1000 cal = 4184 J

    Specific Heat Capacity
    • Specific heat capacity (c) is the amount of heat required to raise the temperature of 1 gram of a substance by 1 degree Celsius (or 1 Kelvin).

    • It is a material property that indicates how much energy is needed to change the temperature of a substance.

      Q = mc\Delta T

      • Where:

        • Q is the heat added or removed

        • m is the mass of the substance

        • c is the specific heat capacity

        • \Delta T is the change in temperature

    Heat Transfer Calculations
    • Heat transfer calculations involve determining the amount of heat exchanged during a process, such as heating, cooling, or phase changes.

    • These calculations often use the specific heat capacity and the temperature change to determine the heat transferred.