Kaplan Gen Chem: Chpts 1-3

Proton

positive charge and mass around 1 amu

Neutron

no charge and mass around 1 amu

Electron

negative charge and negligible mass

What does the nucleus contain

protons and neutrons A

Atomic number

number of protons ma

Mass Number

sum of an element’s protons and neutrons

Atomic Mass vs Atomic Weight

Atomic mass = equal to the mass number, the sum of an element’s protons and neutrons

Atomic weight = weighted average of the naturally occurring isotopes of an element

Isotopes

atoms of a given element that have different mass numbers, differ in number of neutrons

Three isotopes of hydrogen names by order

protium, deuterium, and tritium

Rutherford

first postulated that an atom had a dense, positively charged nucleus that made up only a small fraction of the olume of the atom

Bohr Model of the atom

dense, positively charged nucleus is surrounded by electrons revolving around the nucleus in orbits with distinct energy levels

Quantum definition

energy difference between energy levels, first described by Planck

Quantization

there is not an infinite range of energy levels available to an electron

electron can exist only at certain energy levels

energy increases the farther it is from the nucleus

Atomic absorption spectrum

electron must absorb an amount of energy equal to the energy difference between the two levels

Atomic emission spectrum

the amount of energy emitted is exactly equal to the energy difference between the two levels

Quantum mechanical model

electrons do not travel in defined orbits but rather localized orbits

What is an orbital?

region of space around the nucleus defined by the probability of finding an electro in that region of space

Heisenberg uncertainty principle

it is impossible to know both an electron’s position and its momentum exactly at the same time

Principal quantum number (n)

describes the average energy of a shell

azimuthal quantum number (l)

describes the subshells within a given principle energy levels (s, p, d, f)

magnetic quantum number (m_l)

specifies the particular orbital within a subshell where an electron is likely to be found at a given moment in time

spin quantum number (m_s)

indicates the spin orientation (+ or - 1/2) of an electron in an orbital

electron configuration

spectroscopic notation (combining n and l values as a number and letter) to designate the location of electrons

s = 2

p = 6

d = 10

f = 14

Hund’s Rule

subshells with multiple orbitals fill electrons so that every orbital in a subshell gets one electron before any of them gets a second

Paramagnetic

materials have unpaired electrons that align with magnetic fields, attracting to a magnet

Diamagnetic

materials have all paired electrons, which cannot easily be realigned; they are repelled by magnets

Valence Electrons

electrons in the outermost shell available for interaction (bonding) with other atoms

The Periodic Table of the Elements

organizes the elements according to their atomic numbers and reveals a pattern of similar chemical and physical properties

What do they call rows

periods, based on principal energy level (n)

What do they call columns

groups, elements in the same group have the same valence shell electron configuration

Metals

shiny (lustrous), conduct electricity well, and are malleable and ductile

found on left side and middle of the periodic tab;e

Nonmetals

dull, poor conductors of electricity well, and are malleable

found on right side of the period table

Metalloids

possess characteristics of both metals and nonmetals and are found in a stair-step pattern starting with boron (B)

Effective nuclear charge (Z_eff)

net positive charge experienced by electrons in the valence shell and forms the foundation for all periodic trends

increase from left to right across a period

VE become increasingly separated from the nucleus as the principle energy level (n) increase from top to bottom in a group

Atomic Radius

decreases from left to right across a period and increases from top to bottom in a group

Ionic radius

size of a charged species

cation → generally smaller than their corresponding neutral atom

anion → gennerally larger than their corresponding neutral atom

Ionization energy

amount of energy necessary to remove an electron from the valence shell of a gaseous species; it increase sform left to right across a period and decreases from top to bottom in a group

Electron Affinity

the amount of energy released when a gaseous species gains an electron in its valence shells; it increases from left to right across a period and decreases from top to bottom in a group

Electronegativity

measure of the attractive force of the nucleus for electrons within a bond; it increases from left to right across a period and decreases from top to bottom in a group

Alkali metals

typically take on an oxidation state of +1 and prefer to lose an electron to achieve a noble gas-like configuration; they and the alkaline earth metals are the most reactive of all metals

Alkaline earth metals

take on an oxidation state of +2 and can lose two electrons to achieve noble gas-like configuration

Chalcogens

take on oxidation state of -2 or +6 in order to achieve noble gas configuration.

biologically importatn

halogens

typically take on an oxidation stat of -1 and prefer to gain an electron to achieve noble gas-like configurations; these nonmetals have the highest electronegativities

Noble gases

have a fully filled valence shell in their standard state and prefer not to give up or take on additional electrons; they have very high ionization energies, nonexistent electronegativities and electron affinities

Transition metals

take on multiple oxidation states, explains ability to form colorful complexes with nonmetals in solution and their utility in certain biological systems

Octet rule

elements with be most stbale with eight valence electrons

octet rule exceptions

Elements with an incomplete octet are stable with fewer than eight electrons and include H, He, Li, Be, and B

Elements with an expanded octet are stable with more than eight electrons and include all elements in period electron affinity

Compounds with an odd number of electrons cannot have eight electrons on each element

Ionic bond

The transfer of one or more electrons from an element with relatively low energy to an element with a relatively high electron affinity

Ionic bonding rule

occur between elements with large differences in electronegativity (> 1.7), usually between metals and nonmetals

Cation and Anion

cation = positively charged ion

anion = negatively charged ion

Ionic special features

compounds form crystalline lattices

ionic compounds tend to dissociate in water and other polar solvents

ionic solids tend to have high melting points

covalent bond

formed via the sharing of electrons between two elements of similar electronegativities

As bonder order increases…

bond strength and bond energy increases

bond length decreases

Nonpolar covalent bond

both atoms have exactly the same electronegativity

there is a small different between the attoms < 0.5

Polar covalent bond

there is a significant difference in electronegativities 0.5 to 1.7

not enough to transfer electrons and form an ionic bond

the more electronegative element takes on a partial negative charge

the less electronegative element takes on a partial positive charge

Coordinate covalent bonds

result when a single atom provides both bonding electrons while the other atom does not contribute any

Lewis dot symbols

chemical representation of an atom’s valence electrons

requires a balance of valence, bonding, and nonbonding electrons in molecule or ion

Formal charges

exist when an atom is surrounded by more or fewer valence electrons than it has in its neutral state

Resonance structures

with pi system of electrons

all of the possible configurations of electrons

valence shell electron pair repulsion (VSEPR) theory

predicts the 3D molecular geometry of covalently bonded molecules

electrons arrange themselves to be as far apart as possible from each ohter in 3D space

Electronic geometry vs Molecular geometry

electronic geometry = refers to the position of all electrons in a molecule, whether bonding or nonbonding

molecular geometry = refers to the position of only the bonding pairs of electrons in molecule Po

Polarity of molecules

dependent on the dipole moment of each bond and the sum of the dipole moments in molecular structure

all polar molecules contain polar bonds

nonpolar molecules may contain nonpolar bonds, or polar bonds with dipole moments that cancel each other

Sigma and pi bonds

the describe the patterns of overlap observed when molecular bonds are formed

sigma bonds are the result of head to head overlap

pi bonds are the result of the overlap of two parallel electron cloud densities

Intermolecular forces

electrostatic attractions between molecules

weaker than covalent bonds

London dispersion forces

weakest interactions, but are present in all atoms and molecules

as the size of the atom or structure increases, so does the corresponding London dispersion forces D

Dipole dipole interactions

occur between the oppositely charged ends of polar molecules

evident in the solid and liquid phases but negligible in the gas phase due to the distance between particles

Hydrogen bonds

specialized subset of dipole dipole interactions involved in intra and intermolecular attraction

H-bond is bonded to one of three very electronegative atoms: F, O, N