AP Chemistry

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1. Chemical foundations

2. Atoms, molecules, and ions

3. Stoichiometry

4. Types of chemical reactions and solution stoichiometry

5. Gases

6. Thermochemistry

7. Atomic structure and periodicity

8. Bonding: general concepts

1. Chemical foundations

Terms

scientific method: process of studying natural phenomena, involving observations, forming laws and theories, and testing of theories by experimentation

measurement: observation involving a number and a unit (quantitative)

hypothesis: one or more assumptions put forth to explain the observed behavior of nature

theory: set of assumptions put forth to explain some aspect of the observed behavior of matter

model: see “theory”

natural law: statement that expresses generally observed behavior

law of conservation of mass: mass is neither created nor destroyed

SI system: International System of units based on the metric system and units derived from the metric system

mass: quantity of matter in an object

weight: force exerted on an object by gravity

uncertainty: characteristic that any measurement involves estimates and cannot be exactly reproduced

significant figures: certain digits and the first uncertain digit of a measurement

accuracy: agreement of a particular value with the true value

precision: degree of agreement among several measurements of the same quantity; reproducibility of a measurement

random error: error that has an equal probability of being high or low

systematic error: error that always occurs in the same direction

exponential notation: expresses a number as

⁠

; convenient method for representing a very large or very small number and for easily indicating the number of significant figures

unit factor method: equivalence statement between units used for converting from one unit to another

dimensional analysis: see “unit factor method”

density: property of matter representing the mass per unit volume

matter: the material of the universe

states (of matter): the three different forms in which matter can exist; solid, liquid, and gas

homogenous mixture: mixture with visually indistinguishable parts

heterogeneous mixture: mixture with visually distinguishable parts

solution: see “homogenous mixture”

pure substance: substance with constant composition

physical change: change in the form of a substance, but not in its chemical composition; chemical bonds are not broken

distillation: method for separating the components of a liquid mixture that depends on differences in the ease of vaporization of the components

filtration: method for separating the components of a mixture containing a solid and a liquid

chromatography: the general name for a series of methods for separating mixtures by using a system with a mobile phase and a stationary phase

paper chromatography: method of chromatography that uses a porous paper to separate the mixtures

compound: substance with constant composition that can be broken down into elements by chemical processes

chemical change: the change of substances into other substances through a reorganization of the atoms; chemical reaction

element: substance that cannot be decomposed into simpler substances by chemical or physical means

Concepts

scientific method

make observations

formulate hypotheses

perform experiments

models (theories) are explanations of why nature behaves in a particular way

subject to modification over time

sometimes fail

quantitative observations are called measurements

consist of a number and a unit

involve some uncertainty

uncertainty indicated by use of significant figures

rules to determine significant figures

calculations using significant figures

preferred system: SI system

temperature conversions

⁠

matter can exist in three states

solid

liquid

gas

mixtures can be separated by methods involving only physical changes

distillation

filtration

chromatography

compounds can be decomposed to elements only through chemical changes

2. Atoms, molecules, and ions

Terms

law of conservation of mass: mass is neither created nor destroyed

law of definite proportion: given compound always contains exactly the same proportion of elements by mass

law of multiple proportions: when two electrons form a series of compounds, the ratios of the masses of the second element that combine with one gram of the first element can always be reduced to small whole numbers

atomic masses: the weighted average mass of the atoms in a naturally occurring element

atomic weights: see “atomic masses”

Avogadro’s hypothesis: at the same temperature and pressure, equal volumes of different gases contain the same number of particles

cathode rays: the “rays” emanating from the negative electrode (cathode) in a partially evacuated tube; a stream of electrons

electrons: negatively charged particle that moves around the nucleus of an atom

radioactivity: the spontaneous decomposition of a nucleus to form a different nucleus

nuclear atom: atom having a dense center of positive charge (the nucleus) with electrons moving around the outside

nucleus: the small, dense center of positive charge in an atom

proton: positively charged particle in an atomic nucleus

neutron: particle in the atomic nucleus with mass virtually equal to the proton’s but with no charge

isotopes: atoms of the same element (the same number of protons) with different numbers of neutrons; identical atomic numbers but different mass numbers

atomic number: the number of protons in the nucleus of an atom

mass number: the total number of protons and neutrons in the atomic nucleus of an atom

chemical bond: the force or energy that holds two atoms together in a compound

covalent bond: type of bond in which electrons are shared by atoms

molecule: bonded collection of two or more atoms of the same or different elements

chemical formula: the representation of a molecule in which the symbols for the elements are used to indicate the types of atoms present and subscripts are used to show the relative numbers of atoms

structural formula: the representation of a molecule in which the relative positions of the atoms are shown and the bonds are indicated by lines

space-filling model: model of a molecule showing the relative sizes of the atoms and their relative orientations

ball-and-stick model: molecular model that distorts the sizes of atoms but shows bond relationships clearly

ion: atom or group of atoms that has a net positive or negative charge

cation: positive ion

anion: negative ion

ionic bond: the electrostatic attraction between oppositely charged ions

ionic solid: solid containing cations and anions that dissolves in water to give a solution containing the separated ions, which are mobile and thus free to conduct an electric current

polyatomic ion: ion containing a number of atoms

periodic table: chart showing all the elements arranged in columns with similar chemical properties

metal: element that gives up electrons relatively easily and is lustrous, malleable, and a good conductor of heat and electricity

nonmetal: element not exhibiting metallic characteristics; accepts electrons from a metal

group/family: vertical column of elements having the same valence electron configuration and showing similar properties

alkali metals: Group 1A metal

alkaline earth metals: Group 2A metal

halogens: Group 7A element

noble gases: Group 8A element

period: horizontal rows of the periodic table

binary compounds: two-element compound

binary ionic compounds: two-element ionic compound

oxyanions: anion consisting of an atom of an element and a number of oxygen atoms

acid: substance that produces hydrogen ions in solution; proton donor

Concepts

fundamental laws

conservation of mass

definite proportion

multiple proportions

Dalton’s atomic theory

all elements are composed of atoms

all atoms of a given element are identical

chemical compounds are formed when atoms combine

atoms are not changed in chemical reactions, but they way they are bound together changes

early atomic experiments and models

Thomson model

Millikan experiment

Rutherford experiment

nuclear model

atomic structure

small, dense nucleus contains protons and neutrons

protons: positive charge

neutrons: no charge

electrons reside outside the nucleus in the relatively large remaining atomic vollume

electrons: negative charge; small mass (1/1840 of proton)

isotopes have same atomic number but different mass numbers

atoms combine to form molecules by sharing electrons to form covalent bonds

molecules are described by chemical formulas

chemical formulas show number and type of atoms

structural formula

ball-and-stick model

space-filling model

formation of ions

cation: formed by loss of electron; positive charge

anion: formed by gain of electron; negative charge

ionic bonds: formed by interaction of cations and anions

periodic table organizes elements in order of increasing atomic number

elements with similar properties are in columns (groups)

metals are majority; tend to form cations

nonmetals tend to form anions

compounds are named using system of rules depending on type of compound

binary compounds

type I: contain metal that always forms same cation

type II: contain metal that can form more than one cation

type III: contain two nonmetals

compounds containing polyatomic ion

3. Stoichiometry

Terms

chemical stoichiometry: calculation of the quantities of material consumed and produced in chemical reactions

mass spectrometer: instrument used to determine the relative masses of atoms by the deflection of their ions on a magnetic field

average atomic mass: relative average of atomic masses of an element based on isotopic composition

mole: the number of atoms in exactly 12 grams of pure ¹²C, equal to

⁠

Avogadro’s number: see “mole”

molar mass: the mass in grams of one mole of molecules or formula units of a substance; also called “molecular weight”

conceptual problem solving: method of solving problems involving understanding the concept rather than memorizing the process

mass percent: the percent by mass of a component of a mixture or of a given element in a compound

empirical formula: the simplest whole-number ratio of atoms in a compound

molecular formula: the exact formula of a molecule, giving the types of atoms and the number of each type

chemical equation: representation of a chemical reaction showing the relative numbers of reactant and product molecules

reactants: starting substance in a chemical reaction; appears to the left of the arrow in a chemical equation

products: substance resulting from a chemical reaction; shown to the right of the arrow in a chemical equation

balancing a chemical equation: all atoms present in the reactants must be accounted for among the products

mole ratio: the ratio of moles of one substance to mole of another substance in a balanced chemical equation

stoichiometric mixture: mixture that contains the relative amounts of reactants that match the numbers in the balanced equation

limiting reactant: the reactant that is completely consumed when a reaction is run to completion

theoretical yield: the maximum amount of a given product that can be formed when the limiting reactant is completely consumed

percent yield: the actual yield of a product as a percentage of the theoretical yield

Concepts

stoichiometry

deals with amount of substances consumed and/or produced in a chemical reaction

count atoms by measuring mass of sample

relate mass and number of atoms with average atomic mass

mole

amount of carbon atoms in exactly 12 g of pure ¹²C

6.022 × 10²³ units of a substance

mass of 1 mole of an element = atomic mass in grams

molar mass

mass (g) of 1 mole of a compound or element

obtained for a compound by finding the sum of the average masses of its constituent atoms

percent composition

the mass percent of each element in a compound

⁠

empirical formula

simplest whole-number ratio of various types of atoms in a compound

can be obtained from mass percent of elements in compound

molecular formula

formula of constituent molecules

always integer multiple of empirical formula

same for both molecular substances and ionic substances

chemical reactions

reactants are turned into products

atoms are neither created nor destroyed

all atoms present in reactants must be present in products

characteristics of chemical equation

represents chemical reaction

reactants on left side of arrow; products on right side

when balanced, gives relative numbers of reactant and product molecules or ions

stoichiometry calculations

amounts of reactants consumed and products formed can be determined from the balanced chemical equation

limiting reactant is the one consumed first; limits amount of product that can form

yield

theoretical yield is maximum amount that can be produced from given amount of limiting reactant

actual yield (amount of product actually obtained) is always less than theoretical yield

⁠

4. Types of chemical reactions and solution stoichiometry

Terms

aqueous solution: solution in which water is the dissolving medium or solvent

polar molecule: molecule that has a permanent dipole moment

hydration: the interaction between solute particles and water molecules

solubility: the amount of a substance that dissolves in a given volume of solvent at a given temperature

solute: substance dissolved in a liquid to form a solution

solvent: the dissolving medium in a solution

electrical conductivity: the ability to conduct an electric current

strong electrolyte: material that, when dissolved in water, gives a solution that conducts an electric current very efficiently

weak electrolyte: material that, when dissolved in water, gives a solution that conducts only a small electric current

nonelectrolyte: substance that, when dissolved in water, gives a nonconducting solution

acid: substance that produces hydrogen ions in solution; proton donor

strong acid: acid that completely dissociates to produce an H⁺ ion and the conjugate base

strong base: metal hydroxide salt that completely dissociates into ions in water

weak acid: acid that dissociates only slightly in aqueous solution

weak base: base that reacts with water to produce hydroxide ions to only a slight extent in aqueous solution

molarity: moles of solute per volume of solution in liters

standard solution: solution whose concentration is accurately known

dilution: the process of adding solvent to lower the concentration of solute in a solution

precipitation reaction: reaction in which an insoluble substance forms and separates from the solution

precipitate: insoluble substance or solid that forms in a solution

formula equation: equation representing a reaction in solution showing the reactants and products in undissociated form, whether they are strong or weak electrolytes

complete ionic equation: equation that shows all substances that are strong electrolytes as ions

spectator ions: ions present in solution that do not participate directly in a reaction

net ionic equation: equation for a reaction in solution, where strong electrolytes are written as ions, showing only those components that are directly involved in the chemical change

base: substance that produces hydroxide ions in aqueous solution; proton acceptor

neutralization reaction: acid-base reaction

volumetric analysis: process involving titration of one solution with another

titration: technique in which one solution is used to analyze another

stoichiometric (equivalence) point: the point in a titration when enough titrant has been added to react exactly with the substance in solution being titrated

indicator: chemical that changes color and is used to mark the end point of a titration

endpoint: the point in a titration at which the indicator changes color

oxidation-reduction (redox) reaction: reaction in which one or more electrons are transferred

oxidation state: concept that provides a way to keep track of electrons in oxidation-reduction reactions according to certain rules

oxidation: increase in oxidation state; loss of electrons

reduction: decrease in oxidation state; gain of electrons

oxidizing agent (electron acceptor): reactant that accepts electrons from another reactant

reducing agent (electron donor): reactant that donates electrons to another substance to reduce the oxidation state of one of its atoms

Concepts

chemical reactions in solution are very important in everyday life

water is a polar solvent that dissolves many ionic and polar substances

electrolytes

strong electrolyte: 100% dissociated to produce separate ions; strongly conducts an electric current

weak electrolyte: only a small percentage of dissolved molecules produce ions; weakly conducts an electric current

nonelectrolyte: dissolved substance produces no ions; does not conduct an electric current

acids and bases

Arrhenius model

acid: produces H⁺

base: produces OH⁻

Brønsted-Lowry model

acid: proton donor

base: proton acceptor

strong acid: completely dissociates into separated H⁺ and anions

weak acid: dissociates to a slight extent

molarity

one way to describe solution composition

⁠

moles solute = volume of solution (L) × molarity

standard solution: molarity is accurately known

dilution

solvent is added to reduce the molarity

moles of solute after dilution = moles of solute beore dilution

⁠

types of equations that describe solution reactions

formula equation: all reactants and products are written as complete formulas

complete ionic equation: all reactants and products that are strong electrolytes are written as separated ions

net ionic equation: only those compounds that undergo a charge are written ;spectator ions are not included

solubility rules

based on experiment observation

help predict the outcomes of precipitation reactions

important types of solution reactions

acid-base reactions: involve a transfer of H⁺ ions

precipitation reactions: formation of a solid occurs

oxidation-reduction reactions: involve electron transfer

titrations

measures the volume of a standard solution (titrant) needed to react with a substance in solution

stoichiometric (equivalence) point: the point at which the required amount of titrant has been added to exactly react with the substance being analyzed

endpoint: the point at which a chemical indicator changes color

oxidation-reduction reactions

oxidation states are assigned using a set of rules to keep track of electron flow

oxidation: increase in oxidation state (a loss of electrons)

reduction: decrease in oxidation state (a gain of electrons)

oxidizing agent: gains electrons (is reduced)

reducing agent: loses electrons (is oxidized)

equations for oxidation-reduction reactions can be balanced by the oxidation states method

5. Gases

Terms

barometer: device for measuring atmospheric pressure

manometer: device for measuring the pressure of a gas in a container

mm Hg: millimeters of mercury; unit of pressure, also called a torr; 760 mm Hg = 760 torr = 101325 Pa =1 standard atmosphere

torr: see “mm Hg”

standard atmosphere: unit of pressure equal to 760 mm Hg

pascal: the SI unit of pressure; equal to newtons per meter squared

Boyle’s law: the volume of a given sample of gas at constant temperature varies inversely with the pressure

ideal gas: gas that strictly obeys Boyle’s law

Charles’s law: the volume of a given sample of gas at constant pressure is directly proportional to the temperature in kelvins

absolute zero: 0 K

Avogadro’s law: equal volumes of gases at the same temperature and pressure contain the same number of particles

universal gas constant: the combined proportionality constant in the ideal gas law; 0.08206 L × atm/K × mol or 8.3145 J/K × mol

ideal gas law: equation of state for a gas, where the state of the gas is its condition at a given time; expressed by PV = nRT, where P = pressure, V = volume, n = moles of the gas, R = the universal gas constant, T = absolute temperature; expresses behavior approached by real gases at high T and low P

molar volume: the volume of one mole of an ideal gas; equal to 22.42 liters at STP

standard temperature and pressure (STP): the condition 0℃ and 1 atmosphere of pressure

Dalton’s law of partial pressures: for a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone

partial pressures: the independent pressures exerted by different gases in a mixture

mole fraction: the ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture

kinetic molecular theory (KMT): model that assumes that an ideal gas is composed of tiny particles (molecules) in constant motion

root mean square velocity: the square root of the average of the squares of the individual velocities of gas particles

joule: one kilogram meter squared per second squared

diffusion: the mixing of gases

effusion: the passage of gas through a tiny orifice into an evacuated chamber

Graham’s law of effusion: the rate of effusion of a gas is inversely proportional to the square root of the mass of its particles

real gas: gas that does not strictly obey the assumptions of the ideal gas law

van der Waals equation: mathematical expression for describing the behavior of real gases

atmosphere: the mixture of gases that surrounds the earth’s surface

air pollution: contamination of the atmosphere, mainly by the gaseous products of transportation and production of electricity

photochemical smog: air pollution produced by the action of light on oxygen, nitrogen oxides, and unburned fuel from auto exhaust to form ozone and other pollutants

acid rain: result of air pollution by sulfur dioxide

Concepts

state of a gas

can be described completely by specifying its pressure (P), volume (V), temperature (T), and the amount (moles) of gas present (n)

pressure

common units

1 torr = 1 mm Hg

1 atm = 760 torr

SI unit: pascal

1 atm = 101325 Pa

gas laws

discovered by observing the properties of gases

Boyle’s law:

⁠

Charles’s law:

⁠

Avogadro’s law:

⁠

ideal gas law:

⁠

Dalton’s law of partial pressures:

⁠

, where P_n represents the partial pressure of component n in a mixture of gases

kinetic molecular theory (KMT)

model that accounts for ideal gas behavior

postulates of the KMT:

volume of gas particles is zero

no particle interactions

particles are in constant motion, colliding with the container walls to produce pressure

the average kinetic energy of the gas particles is directly proportional to the temperature of the gas in kelvins

gas properties

the particles in any gas sample have a range of velocities

the root mean square (rms) velocity for a gas represents the average of the squares of the particle velocities

⁠

diffusion: the mixing of two or more gases

effusion: the process in which a gas passes through a small hole into an empty chamber

real gas behavior

real gases behave ideally only at high temperatures and low pressures

understanding how the ideal gas equation must be modified to account for real gas behavior helps us understand how gases behave on a molecular level

van der Waals found that to describe real gas behavior we must consider particle interactions and particle volumes

6. Thermochemistry

Terms

energy: the capacity to do work or to cause heat flow

law of conservation of energy: energy can be converted from one form to another but can be neither created nor destroyed

potential energy: energy due to position or composition

kinetic energy:

⁠

energy due to the motion of an object; dependent on the mass of the object and the square of its velocity

heat: energy transferred between two objects due to a temperature difference between them

work: force acting over a distance

pathway: specific conditions dictating the way energy transfer is divided between work and heat

state function (property): a property that is independent of the pathway

system: that part of the universe on which attention is to be focused

surroundings: everything in the universe surrounding a thermodynamic system

exothermic: refers to a reaction where energy (as heat) flows out of the system

endothermic: refers to a reaction where energy (as heat) flows into the system

thermodynamics: the study of energy and its interconversions

first law of thermodynamics: the energy of the universe is constant; same as the law of conservation of energy

internal energy: a property of a system that can be changed by a flow of work, heat, or both;

⁠

, where ΔE is the chnage in the internal energy of the system, q is heat, and w is work

enthalpy: a property of a system

⁠

, where E is the internal energy of the system, P is the pressure of the system, and V is the volume of the system. At constant pressure the change in enthalpy equals the energy flow as heat

calorimeter: device used experimentally to determine the heat associated with a chemical reaction

calorimetry: the science of measuring heat flow

heat capacity: the amount of energy required to raise the temperature of an object by one degree Celsius

specific heat capacity: the energy required to raise the temperature of one gram of a substance by one degree Celsius

molar heat capacity: the energy required to raise the temperature of one mole of a substance by one degree Celsius

constant-pressure calorimetry: pressure remains constant during the calorimetry process

constant-volume calorimetry: volume remains constant during the calorimetry process

Hess’s law: in going from a particular set of reactants to a particular set of products, the enthalpy change is the same whether the reaction takes place in one step or in a series of steps; in summary, enthalpy is a state function

standard enthalpy of formation: the enthalpy change that accompanies the formation of one mole of a compound at 25℃ from its elements, with all substances in their standard states at that temperature

standard state: a reference state for a specific substance defined according to a set of conventional definitions

fossil fuels: coal, petroleum, or natural gas; consists of carbon-based molecules derived from decomposition of once-living organisms

petroleum: thick, dark liquid composed mostly of hydrocarbons

natural gas: fossil fuel composed primarily of methane

coal: remains of plants that were buried and subjected to high pressure and heat over long periods of time

greenhouse effect: a warming effect exerted by the earth’s atmosphere (particularly CO₂ and H₂O) due to thermal energy retained by absorption of infrared radiation

syngas: synthetic gas, a mixture of carbon monoxide and hydrogen, obtained by coal gasification

Concepts

energy

the capacity to do work or to produce heat

is conserved (first law of thermodynamics)

can be converted from one form to another

is a state function

potential energy: stored energy

kinetic energy: energy due to motion

internal energy for a system is the sum of its potential and kinetic energies

internal energy of a system can be changed by work and heat

⁠

work

force applied over a distance

for an expanding/contracting gas

not a state function

⁠

heat

energy flow due to a temperature difference

exothermic: energy as heat flows out of a system

endothermic: energy as heat flows into a system

not a state function

measured for chemical reactions by calorimetry

enthalpy

⁠

is a state function

Hess’s law: the change in enthalpy in going from a given set of reactants to a given set of products is the same whether the process takes place in one step or a series of steps

standard enthalpies of formation (

⁠

) can be used to calculate ΔH for a chemical reaction

⁠

energy use

energy sources from fossil fuels are associated with difficult supply and environmental issues

the greenhouse effect results from release into the atmosphere of gases, including carbon dioxide, that strongly absorb infrared radiation, thus warming the earth

alternative fuels are being sought to replace fossil fuels

hydrogen

syngas from coal

biofuels from plants such as corn and certain seed-producing plants

7. Atomic structure and periodicity

Terms

electromagnetic radiation: radiant energy that exhibits wavelike behavior and travels through space at the speed of light in a vacuum

wavelength: distance between two consecutive peaks or troughs in a wave

frequency: number of waves (cycles) per second that pass a given point in space

Planck’s constant: constant relating the change in energy for a system to the frequency of the electromagnetic radiation absorbed or emitted;

⁠

quantization: concept that energy can occur only in discrete units called quanta

photon: quantum of electromagnetic radiation

photoelectric effect: electrons emit from the surface of a metal when light strikes it

⁠

: Einsein’s equation proposing that energy has mass

E: energy

m: mass

c: speed of light

dual nature of light: statement that light exhibits both wave and particulate properties

diffraction: scattering of light from a regular array of points or lines, producing constructive and destructive interference

diffraction pattern: pattern of bright and dark spots caused by scattered radiation

continuous spectrum: spectrum that exhibits all the wavelengths of visible light

line spectrum: spectrum showing only certain discrete wavelengths

quantum model: electrons move around the nucleus only in certain allowed circular orbits

ground state: lowest possible energy state of an atom or molecule

standing wave: stationary wave as on a string of a musical instrument; in the wave mechanical model, the electron in the hydrogen atom is considered to be a standing wave

wave function: function of the coordinates of an electron’s position in three-dimensional space that describes the properties of the electron

orbital: specific wave function for an electron in an atom; square of this function gives the probability distribution for the electron

quantum (wave) mechanical model: model for the hydrogen atom in which the electron is assumed to behave as a standing wave

Heisenberg uncertainty principle: principle stating that there is a fundamental limitation to how precisely both the position and momentum of a particle can be known at a given time

probability distribution: square of the wave function indicating the probability of finding an electron at a particular point in space

radial probability distribution: graph of total probability of finding the electron in each spherical shell versus the distance from the nucleus

quantum numbers: numbers that describe various properties of an orbital

principal quantum number (n): quantum number relating to the size and energy of an orbital; can have any positive integer value

angular momentum quantum number (ℓ): quantum number relating to the shape of an atomic orbital; can assume any integral value from 0 to n - 1 for each value of n

magnetic quantum number (m_ℓ): quantum number relating to the orientation of an orbital in space relative to the other orbitals with the same ℓ quantum number; can have integral values between ℓ and - ℓ, including zero

subshell: set of orbitals with a given azimuthal quantum number

nodal surface: see “node”

node: area of an orbital having zero electron probability

degenerate orbitals: group of orbitals with the same energy

electron spin: state of an electron that produces one of two oppositely directed magnetic moments

electron spin quantum number (mₛ): quantum number representing one of the two possible values for the electron spin; either + 1/2 or - 1/2

Pauli exclusion principle: in a given atom no two electrons can have the same set of four quantum numbers

polyelectronic atom: atom with more than one electron

aufbau principle: principle stating that as protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen-like orbitals

Hund’s rule: lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli exclusion principle in a particular set of degenerate orbitals, with all unpaired electrons having parallel spins

valence electrons: electrons in the outermost principal quantum level of an atom

core electron: inner electron in an atom; one not in the outermost (valence) principal quantum level

transition metals: several series of elements in which inner orbitals (d or f orbitals) are being filled

lanthanide series: group of 14 elements fallowing lanthanum in the periodic table, in which the 4f orbitals are being filled

actinide series: group of 14 elements following actinium in the periodic table, in which the 5f orbitals are being filled

main-group elements (representative elements): elements in the groups labeled 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A in the periodic table; group number gives the sum of the valence s and p electrons

first ionization energy: energy required to remove the highest-energy electron of an atom

second ionization energy: energy required to remove the second-highest-energy electron of an atom

electron affinity: energy change associated with the addition of an electron to a gaseous atom

atomic radius: half the distance between the nuclei in a molecule consisting of identical atoms

metalloids (semimetals): elements along the division line in the periodic table between metals and nonmetals; exhibit both metallic and nonmetallic properties

Concepts

electromagnetic radiation

characterized by its wavelength (λ), frequency (ν), and speed (c = 2.9979 × 10⁸ m/s)

⁠

can be viewed as a stream of “particles” called photons, each with energy hν, where h is Planck’s constant (

⁠

)

photoelectric effect

when light strikes a metal surface, electrons are emitted

analysis of the kinetic energy and numbers of the emitted electrons led Einstein to suggest that electromagnetic radiation can be viewed as a stream of photons

hydrogen spectrum

emission spectrum of hydrogen shows discrete wavelengths

indicates that hydrogen has discrete energy levels

Bohr model of the hydrogen atom

using the data from the hydrogen spectrum and assuming angular momentum to be quantized, Bohr devised a model in which the electron traveled in circular orbits

although an important pioneering effort, proved to be entirely incorrect

wave (quantum) mechanical model

an electron is described as a standing wave

the square of the wave function (orbital) gives a probability distribution for the electron position

the exact position of the electron is never known (consistent with Heisenberg uncertainty principle: impossible to know accurately both the position and momentum of a particle simultaneously)

probability maps used to define orbital shapes

orbitals characterized by quantum numbers n, ℓ, m_ℓ

electron spin

described by spin quantum number mₛ, which can have values of ±1/2

Pauli exclusion principle: no two electrons in a given atom can have the same set of quantum numbers n, ℓ, m_ℓ, mₛ

only two electrons with opposite spins can occupy a given orbital

periodic table

by populating the orbitals from the wave mechanical model (aufbau principle), the form of the periodic table can be explained

according to the wave mechanical model, atoms in a given group have the same valence (outer) electron configuration

trends in properties (e.g. ionization energies, atomic radii) can be explained in terms of the concepts of nuclear attraction, electron repulsions, shielding, penetration

8. Bonding: general concepts

Terms

bond energy: energy required to break a given chemical bond

ionic bonding: electrostatic attraction between oppositely charged ions

ionic compound: compound that results when a metal reacts with a nonmetal to form a cation and an anion

Coulomb’s law

⁠

E: energy of interaction between a pair of ions (J)

r: distance between the ion centers (nm)

Q₁ and Q₂: numerical ion charges

bond length: distance between the nuclei of two atoms connected by a bond; distance where the total energy of a diatomic molecule is minimal

covalent bonding: type of bonding in which electrons are shared by atoms

polar covalent bond: covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other

electronegativity: tendency of an atom in a molecule to attract shared electrons to itself

dipolar: property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge

dipole moment: see “dipolar”

isoelectronic ions: ions containing the same number of electrons

lattice energy: energy change occurring when separated gaseous ions are packed together to form an ionic solid

single bond: bond in which one pair of electrons is shared by two atoms

double bond: bond in which two pairs of electrons are shared by two atoms

triple bond: bond in which three pairs of electrons are shared by two atoms

localized electron (LE) model: model that assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms

lone pair: electron pair that is localized on a given atom; not involved in bonding

bonding pair: electron pair found in the space between two atoms

Lewis structure: diagram of a molecule showing how the valence electrons are arranged among the atoms in the molecule

duet rule: observation that hydrogen and helium tend to form the most stable molecules when they are surrounded by two electrons (to fill their valence orbitals)

octet rule: observation that atoms of nonmetals tend to form the most stable molecules when they are surrounded by eight electrons (to fill their valence orbitals)

resonance: condition occurring when more than one valid Lewis structure can be written for a particular molecule; the actual electronic structure is not represented by any one of the Lewis structures but by the average of all of them

resonance structure: valid Lewis structure for a molecule with resonance

formal charge: charge assigned to an atom in a molecule or polyatomic ion derived from a specific set of rules

molecular structure: three-dimensional arrangement of atoms in a molecule

valence shell electron-pair repulsion (VSEPR) model: model whose main postulate is that the structure around a given atom in a molecule is determined principally by minimizing electron-pair repulsions

linear structure: molecular structure for two electron pairs with 180-degree bond angles

trigonal planar structure: molecular structure for three electron pairs with 120-degree bond angles

tetrahedral structure: molecular structure for four electron pairs with 109.5-degree bond angles

trigonal pyramid: molecular structure for four electron pairs in which one side is different from the other three

trigonal bipyramid: molecular structure for five electron pairs consisting of two trigonal-based pyramids sharing a common base

octahedral structure: molecular structure for six electron pairs with 90-degree bond angles

square planar structure: molecular structure for four electron pairs with 90-degree bond angles in one plane; not the ideal structure

Concepts

chemical bonds

hold groups of atoms together

occur when a group of atoms can lower its total energy by aggregating

types of chemical bonds

ionic: electrons are transferred to form ions

covalent: equal sharing of electrons

polar covalent: unequal electron sharing

percent ionic character of a bond X—Y

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electronegativity: relative ability of an atom to attract shared electrons

polarity of a bond depends on the electronegativity difference of the bonded atoms

spatial arrangement of polar bonds in a molecule determines whether the molecule has a dipole moment

ionic bonding

an ion has a different size than its parent atom

an anion is larger than its parent ion

a cation is smaller than its parent atom

lattice energy: the change in energy when ions are packed together to form an ionic solid

bond energy

energy necessary to break a covalent bond

increases as the number of shared pairs increases

can be used to estimate the enthalpy change for a chemical reaction

Lewis structures

show how the valence electron pairs are arranged among the atoms in a molecule or polyatomic ion

stable molecules usually contain atoms that have their valence orbitals filled

leads to a duet rule for hydrogen

leads to an octet rule for second-row elements

the atoms of elements in the third row and beyond can exceed the octet rule

several equivalent Lewis structures can be drawn for some molecules, a concept called resonance

when several nonequivalent Lewis structures can be drawn for a molecule, formal charge is often used to choose the most appropriate structure(s)

VSEPR model

based on the idea that electron pairs will be arranged around a central atom in a way that minimizes the electron repulsions

can be used to predict the geometric structure of most molecules

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