BZ

Chemistry - SCH3U - Unit 1 - Atomic Theory & Periodic Trends

Early theories of atomic structure

How models developed

  • Emperial knowledge (observations/data)

  • Theoretical knowledge (ideas from observations)

    • To explain what see

  • Theories

    • Describe observations in term non-observable idea

    • Explain using idea and model

    • Perdict result

    • Simple as possible

    • Are dynamic

      • They change

Democritus (400 B.C.)

  • Matter compose tiny particle

  • Keep splitting matter constantly will reach point where can not keep going

  • Atomos = uncuttable

  • Atoms different size, constant motion, seperate empty space

  • Every matter is made of that particle

    • Cheese = cheese atom have cheese propertiies

Aristotle (350 B.C.)

  • Reject Democritus theory

  • All matter made of Earth, Fire, Air, Water

  • Each basic substance from combination four specific qualities

    • Cold, dry, hot, moist

  • Accepted for ~2000 years

Dalton (1800)

  • 17th-18th centuries brought scientific revolution

    • Emphasis on scientific method and evidence

  • Billiard ball model (atom as solid sphere)

    • All matter composed tiny indivisible particle called atom

    • All atom of same element have identicle properties

    • Atom of different element have different properties

    • Atom of two or more element can combine in constant ration to form new substances

    • In chemical reactions, atom join or seperate but are not destroyed

      • Law of conservation of mass

  • Did not account for how electrical charges were acquired

JJ Thomson (1897)

  • Invented cathode ray tube

  • Propose idea of subatomic particles

  • First to hypothesize electron

  • Electrons, negative charge

  • Suggest negative electron embedded of surface of positive charge sphere

    • Mostly empty space

    • Plum pudding model

    • Rasin bun model

  • Cathode ray tube experiment

    • See if ray has electrical charge

    • Test charge of ray using opposite charged charge plate and magnet

      • Turned away from negative and towards positve

    • Conclude must have negative charge

    • Particle that make up cathode ray tube 1000 time smaller then hydrogen atom

    • Different mental get same cathode ray tube

Nagaoka (1903)

  • Represented atom as positive charge sphere with ring of negative charge electron

  • Saturnian model

Rutherford (1903)

  • Thought Thomson corrrect

    • Wanted test Thomson theory

  • Gold foil experient

    • Shot alpha (x) particle (tiny positively charged, smaller then atom) through thin gold foil

    • If Thomson model correct, particles would pass through

      • As atom mostly empty space

    • Most did but some did not

      • Found atom have nucleus that occasionally cause alpha particle bounce back

        • Something hard inside atom or something same charge as alpha (x) particle

    • Conclude

      • Nuclues positive charge + contain majority atom mass

      • Atom mostly empty space

      • Electron orbit nucleus

        • Float near nucleus

          • No belief of orbital shells yet

  • Protons

Chadwich (1932)

  • Demonstrated nuclei must contain heavy neutral particles as well as proton

    • Called them neutron

Bohr model of atom

  • Built off other scientists to develop planetary model

Limitations of Rutherford model

  • Why electron not go to nucleus

  • What keep them from escaping and traveling out of barrier

  • What keep proton so close together if like charges repel

Bohr model

  • Electron are confined in energy level

  • Fixed distances from nucleus

  • Electrons in same orbit have energy

    • Closer = less energy

    • Further = more energy

  • Electron do not exist between energy level

  • Always occupy lowest possible energy level

    • Can not just skip first shell and fill next

  • Max electron in energy level is 2n²

    • Where n is energy level

Evidence

  • Line spectrum

    • When atom of element heated will release set distinct colour

    • Lines of colour are called line spectrum

    • Each element has own unique line spectrum

    • Each colour has distinct amount energy

      • Red = lowest

      • Purple = highest

  • Connecting evidence to model

    • Electron absorb energy when heated

    • If enough heat added, electron can jump to higher energy level

      • Excited state

    • When energy is lost will fall back down

      • Ground state

    • Energy lost is released as light to give distinct colour

    • Line spectrum represent transition from excited to ground state

      • Higher jump = higher energy

  • NOTE

    • Even though Hydrogen only one electron will not have just one atom

      • Different atom may get different amount energy

Flame test

  • Heat metal or metal ion use flame to cause electron excited state

  • Moving back down to ground state release energy in form visible light

  • Different metal produce unique colour

Atomic mass

  • If atomic mass = proton + neutron and proton and neutron are whole, why is atomic mass a decimal?

Isotopes

  • 2 or more atom of same element with different atomic mass

  • Due to different number of neutron within nuclei

Similarities

  • Same proton and electron

  • Same appearance and chemical properties

    • Different if isolate single isotope

Differences

  • Different neutron

  • Different atomic mass

Isotopic Abundance

  • Sample magnesium is a mixture of its three isotopes

  • Each isotope is fraction of mixture and has own isotopic abundance

  • Isotopic abundance is fixed

    • Every sample of the element has same proportions of isotopes

  • Eg

    • Magnesium

      • 79% Mg-24

      • 10% Mg-25

      • 11% Mg-26

Average atomic mass and isotopic abundance

  • Average atomic mass (on periodic table) is weighted average of all isotope masses

  • Takes into account isotope masses and percent abundances

  • In weighted average calculation isotope with greatest percentage has biggest influence on average atomic mass

Average atomic mass

  • Need to know: Number of isotopes, mass of each isotope, percentage abundance of each isotope

  • Eg

    • Average atomic mass of Carbon is 12.01u which mean

      • Carbon has more than one isotope

      • One of isotope is very close to 12, one is greater

      • Most abundant is close to 12

Calculating

  • AAM(u) = (%ab_1 x mass_1 ) + (%ab_2 x mass_2 ) + . . . (%ab_{n} x mass_{n} )

  • AAM = average atomic mass

  • u = unified atomic mass unit

  • %ab = percent abundance (put into decimal)

  • mass = mass of isotope

  • Eg

    • Boron-10 (10.01u) and Boron-11 (11.01). Average atomic mass: 10.81u. Let x represent the percentage of abundance in boron-10.

      • 10.81 = x(10.01) + (1-x)(11.0)

      • 10.81 = 10.01x + 11.01 - 11.01x

      • x = 0.20

        • Therefore the abundance of Boron-10 is 20% and Boron-11 is 80% (1-x)

Radioisotopes

  • Unstable isotopes

  • Produce nuclear radiation as decay and become radioactive

Nuclear radiation

  • Type of radiation, description, speed, air breakthrough, barrier

    • Alpha, same nucleus as He(2p,2n,+2), slow, few cm, sheet of paper

    • Beta, e^- , fast, few m, 1-2mm of metal

    • Gamma, no mass high energy electromagnetic rod more wave then particle, speed of light, unlimited, 1m lead or concrete

Half-life

  • Time take for one-half originial radioactive atom to decay

  • Vary between isotopes (cs-142 = 5×10^{15} years vs Polonium-216 = 0.16 seconds)

Periodic trends

Forces within atom

  • Properties can be understood by looking at valence electron

    • Forces of attraction with nucleus

    • Forces of repulsion with other electrons

Core charge

  • How to explain force of extraction of nucleus on outer electron

  • cc = # proton - # inner electron (everything but valence)

  • Measure of “effective nuclear charge“ (not important)

  • Eg

    • Sodium

      • 11 protons

      • Electrons per energy shell

        • 2, 8, 1

      • cc = 11-10

      • cc = 1

    • Flourine

      • 9 protons

      • Electrons per energy shell

        • 2, 7

      • cc = 9-2

      • cc = 7

    • Therefore flourine has a stronger force of attraction than sodium

Electron shielding

  • Measure force of repulsion

  • Amount of sheilding depend on # inner shells

    • Number of sheilds is number of full shells

  • More complex shell = more force of repulsion

  • Eg

    • Sodium

      • Electrons per energy shell

        • 2, 8,1

    • Flourine

      • Electrons per energy shell

        • 2, 7

    • Therefore sodium has a greater force of repulsion

Periodic trends - size

  • Atomic radius

    • Increases down a group

      • Force of repulsion increases

        • From adding shells

    • Decreases moving right

      • Force of attraction increases

        • From adding electrons to existing shells

  • Ionic radius

    • Groups 1, 2, 13-18

      • Cations always smaller than original (core charge increases)

      • Anions always larger than original (newly completed electron shell)

  • NOTE

    • If have neutral atom (not an ion), cc is = valence electron

  • Eg

    • Which element, Li, B, or F has smallest radius

      • Lithium

        • 3 protons

        • Electrons per energy shell

          • 2, 1

        • cc = 3-2

        • cc = 1

      • Boron

        • 5 protons

        • Electrons per energy shell

          • 2, 3

        • cc = 5-2

        • cc = 3

      • Fluorine

        • 9 protons

        • Electrons per energy shell

          • 2, 7

        • cc = 9-2

        • cc = 7

      • Therefore F has the smallest radius. The force of repulsion is the same (same number of inner shells) but F has the largest core charge and therefore the largest force of attraction

Energy trends

  • Ionization energy (IE)

    • Energy required to remove electron

    • Down a group, decreases

      • More repulsive so further from nucleus and easily removed

    • Across period left to right, increases

      • Core charge increases so electrons more attracted

    • Talk about in removing one electron at time

      • 1st, 2nd, . . .

    • Eg

      • Rank from lowest to highest: Be, O, S

        • Beryllium

          • 4 protons

          • Electrons per energy shell

            • 2, 2

          • Electron sheild = 1

          • cc = 4-2

          • cc = 2

        • Oxygen

          • 8 protons

          • Electrons per energy shell

            • 2, 6

          • Electron sheild = 1

          • cc = 8-2

          • cc = 6

        • Sulfur

          • 16 protons

          • Electrons per energy shell

            • 2, 8, 6

          • Electron sheild = 2

          • cc = 16-10

          • cc = 6

        • Therefore S → Be → O

          • S has highest force of repulsion

          • O has highest force of attraction

            • Therefore higher ionization

  • Electron affinity (EA)

    • Energy released when electron is gained

    • Decreased down group

      • Higher repulsion forces makes harder to add and lower energy is released

    • Increase across period left to right

      • Higher core charge so higher energy is released

      • Makes atom more stable

    • Eg

      • Lowest electron affinity: Na, Al, F

        • Sodium

          • 11 protons

          • Electrons per energy shell

            • 2, 8, 1

          • Electron sheild = 2

          • cc = 11-10

          • cc = 1

        • Aluminum

          • 13 protons

          • Electrons per energy shell

            • 2, 8,3

          • Electron sheild = 2

          • cc = 13-10

          • cc = 3

        • Fluorine

          • 9 protons

          • Electrons per energy shell

            • 2, 7

          • Electron sheild = 1

          • cc = 9-2

          • cc = 7

        • Therefore Na has lowest electron affinitiy becuase it has highest electron shielding and lowest attractive force

        • Therefore F has highest electron affinity because it has lowest sheilding and highest attractive force

Electron negativity

  • Tendency for atom to pull bonding electron to itself