AP Chem exam review

imported from rhiannon_schaub's quizlet set

atomic number

The same as the number of protons in the nucleus of an element; it is also the same as the number of electrons surrounding the nucleus of an element when it is neutrally charged.

mass number

The sum of an atom's neutrons and protons

isotopes

Atoms of an element with different numbers of neutrons

Avogadro's number

6.022×10²³ particles per one mole

Moles

grams/molar mass

Standard Temperature and Pressure (STP)

Pressure = 1 atm

Temperature = 273 K

Converting from moles to liters

I mole of gas = 22.4 L

Moles and Solutions

Moles = (molarity)(liters of solution)

percent composition (mass percents)

The percent by mass of each element that makes up a compound. It is calculated by dividing the mass of each element or component in a compound by the total molar mass for the substance.

empirical formula

• Represents the simplest ratio of one element to another in a compound

• Start by assuming a 100 g sample

• Convert percentages to grams

• Convert grams into moles

• Divide each mole value by the lowest of the values

• These values become the subscripts

molecular formula

• Determine the molar mass of the empirical formula

• Divide that mass into the molar mass

x = m/e

x = molar mass/ empirical mass

• Multiply all subscripts in the empirical formula by the value of x

Aufbau principle

States that when building up the electron configuration of an atom, electrons are placed in orbitals, subshells, and shells in order of increasing energy.

Pauli Exclusion Principle

States that the two electrons which share an orbital cannot have the same spin. One electron must spin clockwise, and the other must spin counterclockwise.

Hund's Rule

States that when an electron is added to a subshell, it will always occupy an empty orbital if no one is available. Electrons always occupy orbitals singly if possible and pair up only if no empty orbitals are available.

Coulomb's Law

The amount of energy that an electron has depends on its distance from nucleus of an atom. While on the exam, you will not be required to mathematically calculate the amount of energy a given electron has, you should be able to qualitatively apply Coulomb's Law.

Essentially, the greater the charge of the nucleus, the more energy an electron will have.

Quantum Theory

Max Planck figured out that electromagnetic energy is quantized. That is, for a given frequency of radiation (or light), all possible energies are multiples of a certain unit of energy, called a quantum (mathematically, that's E = hv). So, energy changes do not occur smoothly but rather in small but specific steps.

Energy and Electromagnetic Radiation

ΔE = hv = hc/λ

ΔE = energy change

h = Planck's constant, 6.626×10⁻³⁴ J∙s

v = frequency of the radiation

λ = wavelength of the radiation

c = the speed of light, 3.00×10⁸ m/s

Frequency and Wavelength

c = λv

Inversely proportional

c = speed of light in a vacuum (2.998×10⁸ m/s)

λ = wavelength of the radiation

v = frequency of the radiation

ionization energy

The amount of energy necessary to remove an electron from an atom.

photoelectron spectra (PES)

A chart of the amount of ionization energy for all electrons ejected from a nucleus.

The y-axis describes the relative number of electrons that are ejected from a given energy level.

The x-axis shows the binding energy of those electrons.

Gases will most likely act as ideal under what conditions?

High temperature and low pressure

Energy Levels

s-subshell holds two electrons

p-subshell holds six electrons

d-subshell holds 10 electrons

f-subshell holds 14 electrons

electron configuration

The complete description of the energy level and subshell that each electron on an element inhabits

Heisenberg Uncertainty Principle

States that it is impossible to know both the momentum of an electron at a particular instant.

Atomic Radius Trends

Atomic radius decreases across a period

Atomic radius increases down a group

Cations are smaller than their atoms

Ions are larger than their atoms

Ionization energy

The energy required to remove an electron from an atom.

Electronegativity

Refers to how strongly the nucleus of an atom attracts the electrons of other atoms in a bond.

Periodic Trends

Across the periods

• atomic radius decreases

• ionization energy increases

• electronegativity increases

Down the periods

• atomic radius increases

• ionization energy decreases

• electronegativity decreases

Ionic Bonds

An ionic solid is held together by the electrostatic attractions between ions that are next to one another in a lattice structure.

Occurs between a metal and a nonmetal; electrons are not shared, they are given up by one atom and accepted by another.

Substances with ionic bonds are usually solids at room temperature and have very high melting and boiling points.

Ex. NaCl

Factors Affecting Melting Points of Ionic Substances

1. Charge on ions - a greater charge leads to a greater bond energy

Ex. MgO will have a higher melting point than NaCl

2. Size of ions - smaller ions will have greater attraction

Ex. LiF will have a greater melting KBr

Interstitial alloys

Metal atoms with two different radii combine

Ex. In steel, much smaller carbon atoms occupy the interstices of the iron atoms

Substitutional alloy

Forms between atoms of similar radii

Ex. Atoms of zinc are substituted with copper atoms to create an alloy

Covalent bonding

Bonding in which two atoms share electrons. Each atom counts the shared electrons as part of its valence shell to achieve complete outer shells.

The first covalent bond formed between two atoms is called a sigma bond.