chem

Briefing Document: Introduction to Chemistry - Fundamentals

I. Introduction: Science, Chemistry, and the Scientific Method (Chapter 1.1)

• Definition of Science and Chemistry:Science is defined as knowledge gained through experience or experimentation; it's both an activity and the result of that activity. It's specifically empirical knowledge.

• Chemistry is the study of matter and its interactions with other matter and energy.

• Scientific Method: Investigations are guided by theory and earlier experiments.

• Hypothesis, Law, and Theory:A hypothesis is a possible explanation for an event.

• A law is a statement summarizing a large number of observations (it doesn't explain why; it just states what happens.)

• A theory is an explanation of the laws of nature (it explains why the law is observed).

II. Matter and its Properties (Chapter 1.2)

• Definition of Matter: Matter is anything that has mass and occupies space.

• Mass is the quantity of matter in an object.

• Weight is the force of attraction between an object and other objects, influenced by gravity

• Properties of Matter:A property is anything that can be observed or measured about a sample of matter.

• Extensive properties depend on the size of the sample (e.g., mass, volume).

• Intensive properties are independent of the sample size (e.g., density, color, melting point, boiling point).

• Physical vs. Chemical Properties/Changes:Physical properties can be measured without changing the composition of the sample (e.g., mass, density, color, melting point).

• A physical change occurs without changing the composition of the material (e.g., freezing, melting).

• Chemical properties describe the reactivity of a material (e.g., flammability).

• A chemical change involves at least part of the material transforming into a different kind of matter (e.g., burning methane). "Burning methane in air produces carbon dioxide and water"

• Classification of Matter:Matter is divided into substances and mixtures.A substance is chemically the same throughout. These are either elements (cannot be broken down further, like Aluminum or Oxygen) or compounds (can be broken down into elements, like Sodium Chloride (NaCl)).

• A mixture can be separated into simpler materials by physical methods. Mixtures can be either heterogeneous (composition varies, like rocks) or homogeneous (uniform composition throughout, like air, also called a solution.) Alloys are solutions of a metal with another material.

III. Measurements, Uncertainty, and Units (Chapter 1.3 & 1.4)

• Accuracy vs. Precision:Accuracy is the agreement of a set of measurements with the true value.

• Precision is the agreement among repeated measurements of the same quantity. The document shows examples of results of measurements that are both accurate and precise, accurate but not precise, precise but not accurate, or neither.

• Significant Figures:The number of significant figures includes all digits from the first non-zero digit to the last reported digit, including the uncertainty which is at least ±1 in the last reported digit.

• Zeros preceding the first non-zero digit are not significant.

• Zeros are significant if they are after a non-zero digit and after a decimal point.

• Trailing zeros in numbers without a decimal point may or may not be significant. Use scientific notation to avoid ambiguity.

• Some quantities are not limited by significant figures, including counted numbers (there are 112 students in the class), defined numbers (a dozen is 12 eggs), and powers of ten in exponential notation.

• Uncertainty in Calculations:Addition/Subtraction: The answer should have no more decimal places than the least accurate number (the number with the fewest decimal places).

• Multiplication/Division: The answer should have no more significant figures than the least accurate number (the number with the fewest sig figs).

• Mixed Operations: Follow the normal mathematical order of operations to determine significant figures.

• Rounding:If the digit after the last significant figure is less than 5, round down.

• If it's 5, round to the nearest even number

• If it's more than 5, round up.

• SI Units:The Système International (SI) is the standard for measurements in science.

• Base units include meter (m), kilogram (kg), second (s), kelvin (K), mole (mol), ampere (A) and candela (cd).

• Unit Conversion: A unit conversion factor is a fraction in which the numerator and denominator are equal quantities expressed in different units. For example, "The relationship 1 kg = 1000 g generates two unit conversion factors: (1 kg / 1000 g) and (1000 g / 1 kg)"

• Density is defined as mass per unit volume: d = m/V

• Temperature Conversions: Formulas are provided for converting between Fahrenheit, Celsius, and Kelvin scales. Key temperatures for water are provided in each scale (freezing point and boiling point.)

IV. Atoms, Molecules, and Ions (Chapter 2)

• Dalton's Atomic Theory:Matter is composed of atoms. Atoms are the smallest units of an element that have all the properties of that element.

• An element is composed entirely of one type of atom.

• A compound contains atoms of two or more different elements. The relative number of atoms of each element in a compound is always the same.

• Atoms do not change identity in chemical reactions; only the ways in which they are joined change.

• Laws Related to Atomic Theory:Law of Constant Composition: All samples of a pure substance contain the same elements in the same proportions by mass.

• Law of Multiple Proportions: When the same elements form more than one compound, the mass of one element that combines with a fixed mass of a second element are in a ratio of small whole numbers.

• Law of Conservation of Mass: There is no detectable change in mass when a chemical reaction occurs.

• Subatomic Particles:Atoms are made of: electrons (e-), protons (p+), and neutrons (n0).

• Electrons have a negative charge, protons have a positive charge, and neutrons have no charge. Relative mass: electrons are nearly 0, protons and neutrons are ~ 1.

• "J.J. Thomson demonstrated that electrons were negatively charged" and "Robert A. Millikan performed experiments that determined the charge of the electron as –1.602 x 10-19 coulombs."

• Nuclear Model of the Atom (Rutherford): The atom consists of a small nucleus with positive charge and most of the atom's mass and the rest of the space being mostly empty with the electrons. Rutherford proposed that a hydrogen nucleus was a fundamental particle called a proton.

• Atomic Number (Z) and Mass Number (A):Atomic number (Z) is the number of protons in an atom's nucleus. This defines the element.

• Mass number (A) is the sum of protons and neutrons in an atom's nucleus.

• Isotopes: Atoms of the same element with different numbers of neutrons (same Z, different A). Isotopes are identified by the formula AZX where A is mass number, Z is atomic number and X is the one or two letter symbol of the element. Since Z identifies the element, it is often omitted in writing isotopic symbols

• Ions: Charged particles formed when atoms gain or lose electrons.

• Cations are positively charged (formed when atoms lose electrons).

• Anions are negatively charged (formed when atoms gain electrons). Charge is shown as a superscript after the symbol, example Ca2+. A charge of zero is not shown.

• Atomic Mass Unit (u): 1/12 the mass of one 12C atom. Protons and neutrons have masses approximately 1u.

• Mass Spectrometry: Used to measure the masses and abundances of isotopes, allowing the calculation of the weighted average atomic mass of an element. "Atomic mass = fractionA x isotopic massA + fractionB x isotopic massB + …"

V. Periodic Table (Chapter 2.5)

• Organization: Arranges elements with similar properties in columns (groups/families) and rows (periods).

• Classifications of Elements:Metals: Good electrical conductors, typically shiny. On the left side of the periodic table.

• Nonmetals: Typically non-conductors. On the top right of the periodic table.

• Metalloids: Have properties of both metals and nonmetals; located along the staircase on the periodic table.

• Representative Elements: Groups 1A-8A (or 1,2, 13-18).

• Transition Metals: Groups 1B-8B (or 3-12).

• Inner Transition Metals: The two rows at the bottom of the table: lanthanides and actinides.

• Important Groups:Alkali Metals: Group 1A (1) - soft, reactive metals.

• Alkaline Earth Metals: Group 2A (2).

• Halogens: Group 7A (17) - reactive nonmetals, also called salt-formers.

• Noble Gases: Group 8A (18) - stable, mostly inert gases.

VI. Molecules, Molecular Mass, and Ionic Compounds (Chapter 2.6, 2.7, & 2.8)

• Molecules: Combinations of atoms joined strongly, behaving as a single particle.

• Molecular Formulas: Give the number of every type of atom in a molecule.

• Structural formulas show how atoms are connected.

• Diatomic Molecules: Simplest molecules, containing two atoms (homonuclear if same atom, heteronuclear if different).

• If all atoms in a molecule are the same, it's an element.

• If two or more elements form a molecule, it's a molecular compound (typically formed by nonmetals).

• Molecular mass is the sum of atomic masses of all atoms in the molecular formula, expressed in atomic mass units (u).

• Ionic Compounds: Composed of cations and anions forming neutral species, generally from a combination of metals and nonmetals, with each cation surrounded by several anions.

• Ionic Compound Formulas: Empirical formulas that use the smallest whole-number subscripts to express the relative numbers of ions. The charges must balance to zero.

• Predicting Ionic Charges: Group 1A, 2A, 3B, and Al (Group 3A) elements form cations with charges equal to the group number; Nonmetals from Groups 6A, 7A and N in group 5A form anions with charges of 2-, 1- and 3- respectively.

• Polyatomic Ions: Groups of atoms with a net charge that behave as a single particle.

• Formula Mass of Ionic Compounds is the sum of atomic masses of all atoms in the empirical formula.

• Chemical Nomenclature: An organized system for naming compounds:

• Ionic Compounds: Cation name first, then anion name (anion ending changed to "ide"). Metals with multiple oxidation states are shown with Roman numerals, e.g., Iron (II) vs Iron (III).

• Acids are hydrogen cations combined with an anion. If the anion ends in -ide, add hydro- prefix with -ic suffix to the anion, e.g., hydrobromic acid from bromide. If the anion ends in -ate, change it to -ic, e.g., phosphoric acid from phosphate. If the anion ends in -ite, change it to -ous.

• Molecular Compounds: Use numerical prefixes (mono, di, tri, etc.) to indicate the number of atoms. (mono- is commonly omitted for the first element.) First element name, then second element name, modified to end in -ide.

• Organic Compounds: Hydrocarbons containing only carbon and hydrogen, with alkanes being named using the suffix "-ane" with the appropriate prefix. Cycloalkanes are named similarly with the prefix "cyclo-." Alkanes can have alkyl groups as substituents (methyl, ethyl, etc) and halogens (chloro) as substituents.

• Numerical Prefixes in Names Number Prefix: one=mono, two=di, three=tri, four=tetra, five=penta, six=hexa

VII. Physical Properties of Ionic and Molecular Compounds (Chapter 2.9)

• Comparison: Ionic compounds are typically hard, brittle solids with high melting points; molecular compounds usually have lower melting points, and some are liquids or gases at room temperature.

• Dissociation: Ionic compounds dissociate into individual cations and anions when dissolved in water. Example: "NaCl dissociates into Na+ (aq) and Cl- (aq) in water."

• Electrolytes: Substances that form ions in water, allowing electrical conductivity (many ionic compounds).

• Nonelectrolytes: Substances that dissolve in water as neutral molecules and do not conduct electricity (most molecular compounds).

• The text provides examples showing how electrolytes cause light bulbs to light, while nonelectrolytes do not.