MCAT General Chemistry - Atomic Structure

Kaplan Test Prep

Nucleus

Nucleus

centre of atom, holds most mass, protons and neutrons

fundamental unit of charge (e)

magnitude of charge of a proton or electron (e = 1.6 × 10^-19 C)

atomic mass unit (amu)

mass of one proton or neutron, exactly 1/12 the mass of carbon-12

Proton

subatomic particle with charge of +1 e, mass = 1 amu, found in the nucleus

atomic number (Z)

identifies element, number of protons in one atom

neutron

subatomic particle with no charge of , mass > 1 amu, found in the nucleus

mass number (A)

sum of the protons and neutrons in one atom

Isotopes

Atoms that have the same atomic number but different mass numbers; have the same number of protons but varying numbers of neutrons; referred to by the name of element followed by mass number; same atomic number means similar chemical properties

Electrons

subatomic particle with charge of - 1 e, mass = 1/2000 amu (often considered 0), found outside the nucleus

Electron shell

a given distance from the nucleus, corresponding to a particular level of electrical potential energy

Valence electrons

Electrons furthest from the nucleus; strongest interactions from surrounding environment and weakest with nucleus; involved with bonding

Ion

charged atom

Cation

positively charged ion

Anion

negatively charged ion

Hydrogen Isotope Names

protium : Z=1, A=1

deuterium : Z=1, A=2

tritium : Z=1, A=3

Atomic weight

weighted average of naturally-occurring isotopes of an element; represents mass of ‘average’ atom in amu and mass of one mol of element in g.

Mole (mol)

number of things equal to Avogadro’s number

Avogadro’s number (N_A)

6.02 × 10²³

Ernest Futherford

1910 - proved atom has small, dense, positively charged nucleus - gold foil experiment

Max Planck

1899 - first quantum theory, energy emitted as electromagnetic radiation comes in discrete bundles/quanta - blackbody experiments/ultraviolet catastrophe

Quantum (pl. quanta)

discrete amount of energy

Planck relation

relates the energy of a quantum to frequency via a proportionality constant; E=hf=hν=hc/λ

Planck’s constant (h)

proportionality constant, 6.626 × 10⁻³⁴ J · s

frequency (f/ν)

wave frequency of radiation

Niels Bohr

1913 - developed planetary model of atom; quantized angular momentum

angular momentum (L)

a vector quantity that describes the rotary inertia of an object or system; L=mvr=nh/2π

Rydberg equation

E= - R_H/n²= R_H[1/n_i²-1/n_f²]; negative = attractive force towards nucleus; energy of electron increases at increasing n

Rydberg unit of energy

2.18 × 10⁻¹⁸ J/electron

orbit

defined pathway of an electron at a discrete energy level

ground state

state of lowest energy; all electrons are in lowest possible orbitals

excited state

at least one electron is at a higher energy level

Atomic Emission Spectrum

spectrum of frequencies of electromagnetic radiation emitted due to electrons making a transition from a high energy state to a lower energy state; each element has unique set of energy levels

line spectrum

representation of atomic emission spectra, where each line represents light at a specific frequency

Lyman series

hydrogen emission lines from n ≥ 2 to n = 1

Balmer series

hydrogen emission lines from n ≥ 3 to n = 2

Paschen series

hydrogen emission lines from n ≥ 4 to n = 3

Atomic Absorption Spectrum

the fraction of incident radiation absorbed by the material over a range of frequencies of electromagnetic radiation; electrons absorb specific amounts of energy to get excited; equal to emission wavelengths

orbitals

regions of space where electrons are often localised, holds two electrons of opposite spins

Heisenberg Uncertainty Principle

It is impossible to simultaneously determine, with perfect accuracy, the momentum and the position of an electron.

Quantum numbers

numbers that describe electrons in an atome

Pauli Exclusion Principle

No two electrons in a given atom can possess the same set of four quantum numbers.

Principal quantum number (n)

represents energy level/electron shell, any positive integer value, max # electrons in shell - 2n²

azimuthal (angular momentum) quantum number (l)

shape and number of subshells within given shell; integers btwn 0 and n-1, max # electrons in subshell = 4l+2

Spectroscopic notation

shorthand representation of the principal and azimuthal quantum numbers; l=0=s, l=1=p, l=2=d, l=3=f

magnetic quantum number (m_l)

specifies orbital; integers from -l to l, including 0; 2 electrons per orbital

spin quantum number (m_s)

specifies spin orientation, ±½

paired electrons

two electrons that occupy the same orbital and have opposite spins

parallel spin

electrons in different orbitals but same spin number

Electron Configuration

the pattern by which subshells are filled, as well as the number of electrons within each principal energy level and subshell; can be abbreviated by placing the noble gas that precedes the element of interest in brackets

Aufbau (building-up) principle

Electrons fill from lower- to higher-energy subshells

n + l rule

the lower the sum of the values of the first and second quantum numbers, n + l, the lower the energy of the subshell.

Hund’s Rule

within a given subshell, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins; due to electron repulsion; half-filled and fully filled orbitals have lower energies (higher stability) than other states

paramagnetic

Materials composed of atoms with unpaired electrons will orient their spins in alignment with a magnetic field, and the material will thus be weakly attracted to the magnetic field

diamagnetic

Materials consisting of atoms that have only paired electrons will be slightly repelled by a magnetic field