Chem 11 Unit 8: Atomic theory (2-1)

Properties of Metals (5)

• Solid at room temp (Except Hg)

• Shiny/lustrous when freshly cut/polished

• Good conductors of heat + energy

• Generally malleable & ductile

• Forms cations

Properties of Metalloids (5)

• When temp increases

  • Metal → conductivity decreases

  • Metalloids → conductivity increases

• B, Si, Ge, As, Sb, + Te = confirmed metalloids

• Po + As = debateable

  • Po = metal OR metalloid

  • As = Non-metal OR metalloid

Properties of Non-metals (4)

• Usually gases/brittle solids at room temp (except liquid bromine)

• Bad conductors of heat and energy

• Solids are dull to lustrous in appearance and opaque to translucent

• Forms anions

Periodic table trend principles (3)

1. e- in large orbits = farther from the nucleus, less attraction to p+ bc more shielding

2. More p+ = more attraction for e- from nucleus→ effective nuclear charge (Zₑff)

3. e- in the same subshell repel each other

Atomic Radius (3)

Measure of the radius from the nucleus to the Ve- in an atom

• Down group → AR increases: More shells (higher n value) + shielding

• Across period → AR decreases: Zeff & Protons increases (Added e- joins same orbital)

Ionic Radius (4)

AR of an ion

• Cations = smaller than neutral

  • Loses electrons so AR decreases

• Anions = larger than neutral

  • Gains elections so AR increases

Ionization Energy (IE) (6)

Amount of energy needed to remove an e-

• First ionization energy (IE₁): the energy required to remove the first e-

• Trend = opposite to AR, small AR → high IE

  • Elements on left need to get rid of e- to become stable

  • Elements on right need to gain e- to become stable

• Exceptions: O = lower IE₁ than N bc O wants to lose 1e- to become half filled

Electronegativity (EN) (5)

The ability to attract e-

• Measurement of an atom’s Zeff on OTHER atoms

• Atoms w/ high EN pull bonded e- closer to their nuclei

• Noble gasses have no EN value (shells stable)

• EN has the same trend as IE

Metallic Character (2)

• Down group = more metallic

• Across period = less metallic

Intramolecular interactions (2)

• Occur WITHIN molecules/compounds/crystals

• Includes: ionic, covalent, metallic bonds

Ionic Bonds (5)

• Forms when large difference between EN and IE of both atoms (EN >1.8)

• Metal w/ low IE/EN transfers Ve- → Non-metal w/ high IE/EN

• Electrostatic attraction-both atoms become ions after transfer

• Sometimes forms into a crystal lattice (Ex. NaCl)

  • Makes compound have a high melting point (Caused by a vast number of attractive forces in structure)

Covalent bonds (gen) (3)

• Non-metal and non-metal

• Share valence electrons

• Tend to have lower MP than ionic compounds

Polar Covalent Bonds (3)

• Forms when EN = 0.5-1.8

  • When e- are pulled more towards one atom bc of higher EN, one side of the moc becomes more negative than the other side (unequal sharing of e-)

    • Causes dipole (One side more negative)

Non-polar Covalent Bonds (7)

• Forms when EN < 0.4

  • Between identical/same charged non-metals

    • EN value = same; ΔEN = 0

  • Collision of atoms = electron clouds overlap

    • Attractive forces (nucleus) may begint o overlap repulsive forces (e-)

      • Pair of Ve- shared between both

      • The electrons will not be more attracted to one nuclei over the other (equidistant)

Metallic Bonds (3)

• Protons swarmed by electrons

  • “sea” of electrons makes metal malleable + conductive

  • Forms metallic crystals

Early Atomic Theory Timeline (6 ppl)

Democritus (he was right all along)

• First coherent atomic theory

  • Matter is made of tiny indivisible particles called atoms

Aristotle (More influential)

• Everything is made up of 4 elements: Fire, air, earth, water

• Criticized Democritus because it didn’t explain properties of matter or chemical reactions

John Dalton

• First “modern” atomic theory (Revived Democritus’ theory through experiments)

  • Atoms = solid spheres (different types of spheres make up different elements)

• Proposed Billiard Ball Model

  • Spheres arranged in a triangle

J.J. Thomson

• Discovered electron (Atoms are not the smallest particles of matter - there are more inside)

  • Aluminum Plates in a Cathode Ray Tube

• Proposed Plum Pudding Model

  • Positive sphere w/ negative charged particles

Ernest Rutherford

• Discovered nucleus (proton confirmed + neutron proposed)

  • Gold Foil Experiment: A lot of α particles went through, some rebounded. Like a net.

• Proposed Planetary/nuclear Model

  • Positive charge in center/nucleus w/ electron surrounding it

  • Flaw: What is stopping the e- from meeting the p+? Atoms would just collapse.

Niels Bohr

• Positive nucleus w/ electrons orbiting in shells of different energy levels (higher shell = higher energy)

  • Chadwick proposed neutrons

Periodic Table history (4 ppl - not important?)

William Odling:

• Proposed elements could be divided into 13 groups based on chemical & physical properties

John Newlands:

• Assigned H a mass of 1

• Law of Octaves: Every 8th element shares similar properties (only up to Ca)

  • Did not allow for prediction of new elements. Table had to be rearranged each time a new element was discovered.

Dimitri Mendeleev:

• Organized elements by atomic mass and properties.

  Recognized periodic recurrence of properties, leaving gaps for undiscovered elements.

• His (first widely recognized) table helped chemists organize data and predict new properties.

• Periodic Law: Elements ordered by increasing atomic mass (later revised to atomic number) exhibit similar recurring properties.

Henry Moseley:

• Proposed atomic number - revised periodic law slightly

• Elements still grouped by properties, similar properties in the same column

• Added a column of elements Mendeleev didn’t know about

Modern table:

• Strutt & Dorn added noble gases in 1894

• Lanthanides & Actinides added mid 1900’s

• Continuously changing, currently at 118 elements

Energy Levels & Orbitals (4)

• Energy difference betw 2 levels is called the QUANTUM energy

• Electrons occupy regions of space called orbitals in a particular energy level (THEY DON’T ORBIT)

• SHELL: The set of orbitals having the same n value

  • Ex. 3rd shell = 3s, 3p, 3d

• SUBSHELL: A set of orbitals of the same type

  • There’s a set of five 3d orbitals in the 3rd subshell

Core Notation/Exceptions (4)

• Anions: Add an electron to the last unfilled subshell

• Cations: Remove an electron from the highest energy subshell

  • Ex. 4s before 3d

  • Ex. np e- before ns e- before nd e-

• Exceptions: Cr & Cu - d subshell wants to be half or fully filled (most subshells do, but esp d)

• Valence: all e- except core e- and filled d or f subshells. Noble gases and isoelectronic atoms/ions have 0 valence.

Orbital rules (3)

• Aufbau Principle: Electrons fill the LOWEST energy orbitals first.

• Pauli Exclusion Principle: Max of 2 electrons per orbital

• Hund’s Rule: Each orbital must have a single electron before doubling up with an e- of opposite spin

Isotopes & Percent abundance (3)

• Isotopes: Different form of same element - Same # protons, diff # neutrons

  • Mass # is always WHOLE (Proton + Neutron)

• Atomic mass = weighted avg of all existing isotopes

Emission Spectra (7)

• When light hits electrons, they jump up a shell

  • If it isn’t the right coloured light, it won’t move

  • Ex. e- that needs red light won’t move if it’s hit with a yellow light

• When they drop down a shell they release specifically coloured light

• Proves the existence of e- shells

  • Solid lines of colour are observed

  • If there was no shells, it’d be a blur of colour

Families / Groups (10)

G1: Alkali Metals

• Very reactive

• Reactivity decreases down group

• Forms 1+ ions

G2: Alkali Earth Metals

• Reactive but not as much as Alkali Metals

• Reactivity decreases down group

• Forms 2+ ions

G3-12: Transition Metals

• D-orbitals filled last (except Cr, Cu, + Au)

• Forms many ions (multivalent)

G13: Boron Group

• Forms 3+ ions

G14: Carbon Group

• Forms 4+ ions (C can form 4- ions)

G15: Nitrogen Group

• Forms 3- ions

G16: Oxygen Groups

• Forms 2- ions

G17: Halogens

• Very reactive

• Reactivity increases up group

G18: Noble Gases

• Very low reactivity, very stable, full Ve- shell

Hydrogen

• Shares properties w/ many families

Octet Rule (5)

• All atoms form stable octet when bonding

  • Stable + low-energy config.

• EXCEPTIONS (sometimes):

  • H makes 1 bond

  • B takes less

  • S and P take more

Intermolecular Forces (4)

• Occurs BETWEEN covalent molecules

• Strength/weakness affects state of matter + chemical properties

  • Strong IMF = High BP (solid/liquid)

  • Weak IMF = Low BP (liquid/gas)

London Dispersion Forces (11)

• Weakest force

• Between 2 non-polar molecules

• Always present

• Creates temporary dipole

  • E- from molec1 → nucleus of molec2

    • Molec2 e- repulsed by molec1 e-

• Attraction increases w/ atomic number

  • Dipole increases when more e- is added

• Attraction increases w/ larger molecules

  • More spaces to attract

• If only LDF present → low BP + MP

Dipole-dipole (4)

• 2nd strongest force

• Between 2 polar molecules

• δ+ from molec1 attracted to δ- from molec2

• Stronger bonding = Higher BP + MP

Hydrogen Bonding (5)

• Strongest force

• Type of dipole-dipole (only for H)

• H bonded with atoms with high EN

  • Ex. N, O, F

• Molecules share a H