IB CHEM SL midterm review (units 1-7)

Pure substance

Pure substance

Chemically bonded with definite and constant composition

Compound

Molecules or lattices with fixed ratio. Have DIFFERENT properties from their elements (but same # protons)

Mixture

Variable, not constant composition, physically bonded. Each substance in a mixture retains its properties (like boil pts)

Homogenous Mixture

All substances in same “phase” w/ no visible boundary

Heterogeneous Mixture

2+ “phases” with a visible boundary

Chemical Properties

Properties in composition that change due to a chemical reaction (flammability, combustion, oxidation, acidity)

Physical Properties

Observed/Measured in a reaction (luster, density, mass, color, size, melt/boil pts, smell, texture, conductivity, electromagnetivity)

Empirical Formula

Simplest whole number ratio of ATOMS in a compiund

Linear P=T

P v T?

linear V=T

V v T

Inverse PV=1

P v V

Linear P=1/V

P v 1/V

Proton

Mass=1 Charge=+1

Neutron

Mass=1 Charge=neutral

Electron

Mass=negligible Charge=-1

A

Mass number (pro+neu)

Z

Atomic number (pro)

X

Element

n

Ions/charge

Isotope

Same element (protons) but different mass (neutrons)

similar chemical but different physical properties

RAM (relative atomic mass)

Weighted average of natural isotope abundances for an element

Carbon dating

14C

Radiotherapy

60Co

Diagnose diseases (medical tracers)

131 and 125 I

Released

Energy is _____ when electrons fall energy levels

Absorbed

Energy is ___ when electrons rise levels

Electromagnetic radiation / light

Y-rays, X-rays, UV, vis light, IR, microwave, radiowave

UV/ultraviolet

Big jumps/falls to/from n=1

Visible light

Medium jumps/falls to/from n=2

Infrared/IR

Small jumps/falls to/from n=3

Discrete e-s

Each element produces a line spectrum instead of a continuous spectrum because ____ exist in discrete levels. Higher levels converge at high energy, excited elements emit a light spectrum

Rows or period

You can tell an elements outermost main energy level by looking at what on the periodic table

Columns or groups

You can tell an elements outermost sub level by looking at what on the periodic table

Spin

Electrons in the same orbital must have different ____.

S orbital

Px orbital

Py orbital

Pz orbital

Aufbau

This principle says to fill the lowest energy level first (this is why 4s fills before 3d!)

Hund

This rule says that in p and d sub levels, fill all orbitals singularly before doubling up.

REMEMBER Cr and Cu steal a 4s electron into their 3d!!!

Transition metals

____ loses electrons from 4s first when they become cations.

3 biggest factors affecting periodic trends

1. # of protons: ENF’s are protons pulling power on val electrons. Added protons across the period pull electrons closer.

2. Outermost energy level: adding levels = bigger atoms

3. E-E repulsion: negative electrons repel each other

Atomic radius

__ increases down a GROUP because energy levels are added.

__ decreases across a PERIOD because more protons pull the same core electrons

Ionic radius

__ increase down a GROUP b/c more energy levels

__ decrease across a PERIOD because more protons pulling electrons

Increases

Across a period, the melting point of metals ___ because atomic radius decreases, and more valence electrons and more IMFs are holding atoms together

Decreases

The melting point of metals ___ down a group because atomic radius increases and there are less IMF’s and valence electrons.

Increases

Down a group, the boiling point of halogens ___ because halogens are diatomic molecules. There are more electrons to completely fill the outer shell with valence electrons

Diatomic elements

Br I N Cl H O F

Ionization energy

__ is the energy required to remove one mole of electrons from one mole of gaseous atoms

K(g) —> K+(g) + e-

Decrease, increase

ionization energy/electronegativity/ electron affinity all ___ down a group, and ___ across a period because smaller atomic radii and they have more pull on electrons

Electronegativity

An atoms attraction of a shared pair of e-s

Electron affinity

The energy change when an electron is added to an isolated gaseous atom.

Cl(g) + e- —> Cl-(g)

Basic

Na2O(s) + H2O(l) —> 2NaOH(aq)

MgO(s) + H2O(l) —> Mg(OH)2(s)

Acidic

SO3(g) + H2O(l) —> H2SO4(aq)

P4H10(s) + H2O —> 4H3PO4(aq)

Amphoteric

Term for when a compound like aluminum oxide can act as either an acid or a base.

Completely steal

Ionic compounds are formed when atoms in the compound ___ electrons.

1.8

Elements forming ionic bonds should have an electronegativity difference of at least _

Ionic bonding

Electrostatic attraction between positive cations and negative anions (typically between metals and nonmetals) in a lattic

Ionic compound

Very high melt and boil point

No electrical conductivity in solids but there is in molten liquids and aqueous solutions

Soluble in polar solvents with weaker bonding

Insoluble in nonpolar solvents

Evenly share

covalent bonds form when atoms in a compound __ electrons

Covalent bonding

Electrostatic attraction between 2 nuclei and a share pair of electrons

0-1.8

Elements forming covalent bonds should have an electronegativity difference of __

Nonpolar covalent bond

0-0.5

Polar covalent bond

0.5-1.8

Increases

As electron domains decrease, bond angles …

Decrease

As unbonded electron pairs increase, bond angles

Resonance structure

Multiple Lewis structures are possible

average of bonds

The “true structure” of resonating bonds are an …

London dispersion forces

Also called temporary dipoles or induced dipoles. Non polar bonds that TEMPORARILY become polar between non polar molecules. Essentially exist between all molecules

Stronger

More linear molecules have more surface area and therefore have ___ temporary dipoles

stronger

Molecules with larger masses form __ temporary dipoles

Dipole dipole forces

Between permanently polar molecules and are stronger than dispersion forces

Hydrogen bonds

Extra strong polarities between molecules. H is attached to N/O/F, and that is bonded to another non bonding e pair on another N/O/F

MUCH stronger than dipole dipole bonds

Giant covalent structure

Infinitely repeatable structure such as diamond.

Not considered molecules because the number of atoms are VARIED- not set like in molecules.

have a very high melting point

Allotrope

Multiple forms in which an element can exist.

Ex. Graphite graphene carbon nanotube (can all conduct electricity)

Ex. O2, O3, O4, O8 (also conduct)

Low

Element forming metallic bonds tend to have __ electronegativities.

Metallic bonding

Electrostatic attraction between cation lattice and “sea” of delocalised electrons (not attached to any single atoms!)

High

Do metals have high or low melting point

Smaller cationic radii, higher ionic charges (more electrons)

Two factors that increase metallic bond strength

Delocalised electrons can move

Solid metals can conduct electricity because …

Cations can slide past each other

Metals are malleable because…

Alloy

Homogenous mixture of 2+ metals (steel bronze brass)

___ are harder/less malleable because different size cations disrupt the lattice and cations can no longer slide.

also have lower boil point because different side atoms make arrangement less regular than a pure metal and bonds are weaker

Temperature

Avg kinetic energy in C or K

Heat

Energy in J or kJ

Heat capacity

How much energy needed to raise 1g substance by 1°C

Resist

Higher specific heats mean substances ___ temp changes

Calorimeter

Device that insulates system (reaction) and surroundings (water) from outside world. Measures temp

Calorimetry assumptions

1. No heat lost to outside surroundings

2. Solutions ONLY water

3. Complete combustion occurred

Chemical bonds

Where energy is stored within a molecule

Break

Energy is required to ___ bonds

Form

Energy is released to ___ bonds

enthalpy change

Amount of energy gained or lost during reaction

stronger

Stable molecules contain __ bonds

Weaker

Molecules with high potential energy contain ___ bonds

Exothermic

Endothermic

Release

Exothermic reactions __ heat

Absorb

Endothermic reactions ___ heat

weaker

Reactants in Exothermic reactions have __ bonds than products

Bond enthalpies

Averages across multiple compounds (in data book). Less specific than experimental calorimetery enthalpy etc