Honors Chemistry Midterm Key Terms and Concepts

cm to inches

2.54 cm = 1 in

giga (G)

G 10^9

mega (M)

M 10^6

kilo- (k)

k 10^3

hecto- (h)

10^2

deka- (da)

10

deci- (d)

0.1

centi (c)

10^-2

milli- (m)

10^-3

micro- (μ)

10^-6

nano- (n)

10^-9

solid

Low energy

Defined shape

Defined volume

Ordered

liquid

Move freely

Takes shape of container

Defined volume

Medium energy

*** aqueous (dissolved in water)

gas

Undefined shape

Undefined volume

Always fills the container/space

High energy

Physical Properties

Substance fundamentally same

reversible

Examples:

Density

Melting

Mixing

Dissolving

Boiling

Chemical Properties

Having to do with a reaction

Substance has fundamentally changed

Irreversible

Examples:

Electrolysis

Rust

Fission

Burning

Density Equation

density = mass/volume

density of water

1 g/mL

less dense

floats

more dense

sinks

elements

Only one type of atom

Represented by an atom

Examples:

O2

N2

Cu

Na

compounds

A combination of different atoms/elements

Represented by a molecule

Examples:

CO2

H2O

NaCl

C6H12O6

ionic compounds

Give/takes an electron

Transfer electrons

Metal + Nonmetal

Conduct electricity in aqueous form

Examples:

NaCl

UF6

covalent compounds

Shares an electron

Nonmetal + Nonmetal

Examples:

CO2 HClO

CH4

H2O

HCl

NH3

organic compounds

ALL COVALENT BONDS

Natural, has to do with living things

CHONPS

Examples:

C6H12O6

CH4

mixtures

at least two substances

separate physically

Examples:

air

saltwater

pure substances

one substance

separate chemically

Examples:

O2

N2

CO2

H2O

heterogenous mixture

different composition

example: ocean water

homogenous

same distribution (uniform)

well mixed

example: air

solutions

homogeneous mixture in aqueous form

examples: milk, soda, alcohol

alloy

homogeneous mixture of metals

examples: bronze, brass, steel, sterling silver, white gold, rose gold

physical methods to separate homogeneous or heterogeneous mixtures

Distillation - separation based on boiling point

handpicking

magnetism - rone, nickel, cobalt

density - floating/sinking

filtration - separates solids and liquids

sieving - separates solids by size

evaporation

decanting - pouring off the top

isotopes

atoms with the same number of protons but different numbers of neutrons

atomic number

number of protons

mass number

number of protons + number of neutrons

cation

electrons are LOST, leaving a positively-charged ion

anion

electrons are gained, leaving a negatively-charged ion

average atomic mass

weighted average of the masses of an element's isotopes

two factors: percent abundance, mass number

alpha particle

helium atom

can be stopped by paper, skin, air

beta particle

electron

can be stopped by clothing, plastic

gamma ray

0/0 y - in the form of energy

can be stopped by lead, concrete

neutron

1/0 n

can be stopped by water, concrete, chemicals

proton

1/1 p or 1/1 H

half-life

time required for half of the original sample to decay

always the same within an element, but times vary dramatically

fusion

combining two lightweight nuclei to form something heavier

fission

splitting a heavy nucleus into two with smaller mass numbers

electromagnetic spectrum

gamma, x-rays, ultraviolet, infrared, microwaves, radiowaves (FM, shortwave, AM)

blue - high frequency/high energy

red - long wavelengths

Photon Energy (Planck Equation)

c=(lambda)(nu)

E = photon energy

h= Planck constant = 6.6261 x 10^-34 J*S

c = speed of light = 3 x 10^8 m/s

lambda = photon wavelength

nu = photon frequency

orbital labels

-The number tells the principal energy level.

-The letter tells the shape.

-The letter s means a spherical orbital.

-The letter p means a two-lobed orbital. The x, y, or z subscript on a p orbital label tells along which of the coordinate axes the two lobes lie.

orbitals different shapes

s, p, d, f

orbital filling rules

Aufbau Principle - fill from lowest energy to highest

N = 1 → 2 → 3 → etc.

s → p → d → f

Hund's Rule - orbitals fill with one electron before doubling up (think about a school bus)

Pauli Exclusion Principle - no two electrons can be identical, so the two that share an orbital have opposite spins

Democritus (400 BCE)

-all things are made of atoms, the fundamental invisible particles

-atoms can't be destroyed

-atoms are separated by the void, or empty space

-atoms come in different shapes and sizes, which explains their various properties

Dalton (1803)

-elements are composed of tiny particle called atoms

-atoms of a given element are identical

-atoms of different elements are different

-atoms of one element can combien with other elements to form compounds

-atoms are individsible in chemical processes, which involve the rearrangement of atoms

Thompson (1897)

Plum pudding model

-positively-charged sphere with negatively-charged electrons embedded

-stated that the atom was a single, uniform sphere of positive charge with negatively-charged electrons scattered throughout

Millikan (1909)

oil drop experiment

-proved that electrons had a fixed, negative charge that didn't vary

-discovered the charge on a single electron, known as the elementary charge (e)

Rutherford (1911)

Gold foil experiment

-showed that the positively-charged nucleus was localized over a very tiny volume of an atom, but took up most of its mass

-atoms are mostly made of empty space

-electrons orbited the nucleus in the center of an atom

Bohr (1913)

-electrons orbit the nucleus in specific paths or energy levels

-each shell has a specific max # of electrons it can hold

-tabeled using the principal quantam # (n) or the K,L,M,N,O,P,Q lettering system

-electrons can absorb energy to jump to higher levels or emit energy when returning to lower energy levels

Schrodinger (1926)

-describes electrons not as particles following fixed paths, but as waves existing in specific regions called orbitals

-model focuses on the probability of finding an electron in a specific region

- location of the electrons is referred to as an electron cloud

-varied densitives: a high density where an electron is most likely to be and a low density where an electron is least likely to be

Chadwick (1932)

-discovered the neutron, revealing that the atomic nucleus contains both positively-charged protons and neutral neutrons

-contributed to the discovery of nuclear fission and the development of the atomic bomb

-measured that the mass of a neutron was about the same as a proton

one mole

contains Avogadro's number of units

atomic mass on table is how many grams of that element are in one mole

Avogadro's number

6.022 * 10^23

writing formula from percentages

Assume 100 g (if necessary)

Determine how many moles of each element there are

Divide by the smallest number of moles to get whole number ratios

Multiple by 2 or 3 until every number is whole

molecular formulas

Compare empirical formula mass to the molar mass. They should be whole number multiples off.

Molecular formula = (empirical formula)n

Molar mass = n * empirical formula mass

diatomic elements

H N F O I Cl Br

charges of some metals

Ag - 1+

Zn - 2+

Cd - 2+

Ga - 3+

In - 3+

covalent compounds - prefix of second element

1 mono

2 di

3 tri

4 tetra

5 penta

6 hexa

7 hepta

8 octa

9 nona

10 deca

electronegativity

the ability of an atom to attract electrons when the atom is in a compound

ionization energy

The amount of energy required to remove an electron from an atom

atomic radius

one-half the distance between the nuclei of identical atoms that are bonded together

types of bonds

polar covalent, nonpolar covalent, ionic, metallic

dipole moment

Tail at positive end, arrow pointing toward negative end.

Higher electronegativity = partially negative

Lower electronegativity = partially positive

octet rule

all elements (except for H and He) want to reach 8 electrons in their outer shell. This is the most stable form. H and He are content with 2 (they're small)

lone pair

pair of electrons that aren't used in bonding

lewis structure exceptions

B and Be - smaller elements that often have less than 8 electrons

Larger noble gases and halogens - can have 'expanded octets' and hold more than 8 electrons. Often S, P, Xe, I, Cl

Atoms with odd numbers of electrons will have a single unpaired. This is called a free radical - often seen in N and O containing compounds.

types of intermolecular forces

1) London dispersion forces

2) Dipole-dipole interactions

3) Hydrogen bonds

4) Anything with an ion

Weakest : London - Dipole - Hydrogen : Strongest

intermolecular forces

forces of attraction between molecules

LDF

Electrons exist in a cloud

The cloud has a certain amount of randomness associated with it

At certain moments, there is a probability that all electrons are on the same side of the atom, creating an "instantaneous dipole"

Instantaneous dipole is a minor force of attraction

More electrons = more momentary dipoles = larger molecules

Dipole-Dipole

Polar molecules have dipoles

When multiple polar molecules are put together, the oppositely charged ends are attracted to each other

Stronger than an instantaneous dipole

More polar, less distance = stronger force of attraction

Hydrogen Bonding

Specific type of dipole-dipole

Extra strong because Hydrogen is tiny and happen when it's bound to something SUPER electronegative

Valid for H-F, H-N, and H-O

Water + H-Bonds

Capillary action

Ice crystalline structure

Glaciers, ice floating

Viscosity

Surface Tension

High specific heat

Effects of IMFs

Boiling point - stronger forces = higher BP

Melting point - stronger forces = higher MP

Vapor pressure - stronger forces = lower VP

Pressure of gas ABOVE the liquid

Ex: acetone (nail polish remover) has a really high vapor pressure

Viscosity - stronger forces = higher viscosity

Surface tension - stronger forces = stronger surface tension

Ideal Gas Law Assumptions

1) nonpolar

2) small

3) high temperature

4) low pressure

Kinetic Molecular Theory

1) continuous random motion

2) particles have no volume

3) pressure = collisions w/ wall

4) no attractions or repulsions

5) kinetic energy = temperature

Ideal Gas Law

PV = nRT

Boyle's Law

Charles' Law

Avogadro's Law

Amonton's Law

Combined Law

Dalton's Law of Partial Pressures