11 - SCH3U1- Chemistry

see Prev Chem Notes

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Notes

Atomic Theory Notes

Key Scientists and their Contributions

• Leucippus & Democritus: First proposed the idea of matter being composed of indivisible particles called “atomos.”
  • Atmos: Cannot be cut/split
• John Dalton: Considered the father of modern atomic theory, proposed that atoms are indivisible and indestructible particles.
• J.J. Thomson:
  • Discovered the electron.
  • Proposed the Plum pudding model where electrons are embedded in a positively charged sphere.
• Niels Bohr:
  • Introduced the concept of quantized energy levels in atoms.
  • Proposed the Bohr model where electrons orbit the nucleus in fixed paths or shells at different energy levels.
• Ernest Rutherford:
  • Conducted the gold foil experiment.
  • Discovered the nucleus at the center of the atom.
• James Chadwick
  • Discovered the neutron along with Ernest Rutherford

History / Progression

• Asirstotle
  • Proposed that all matter is composed of:
    • Earth
    • Air
    • Water
    • Fire
      John Dalton (1766-1844)
  • All mater is made of tiny particles called atoms. They cannot be changed but are divisible.
  • Atoms of one element cannot be converted to any other
  • All the atoms of one element have the same properties
  • From his work the following was found
    Lavoiser: Law of conservation of mass
    Proust: Law of constant composition
    All atoms have a particular combing capacity
• James Chackwidck (1932)
  • Bombarded $Be$ with
• Neils Bohr (1885-1962)
  • Electrons exist in circular orbits (planetary) with electrostatic force from the nucleus holding them in place.
    • Electrons can exist only in a series of allowed orbits (energy levels). therefore the energy of electrons can be quantized
  • An electronic in a stable orbit does not radiate energy
  • Electrons can jump between energy levels by absorbing and emitting photos carrying an amount of energy equal to the jump

Rutherford’s Gold Foil Experiment

• Experiment: Alpha particles were fired at a thin gold foil.
• Observations:
  • Most particles passed through undeflected.
  • Some particles were deflected at large angles.
  • A few particles bounced back.
• Conclusions:
  • The atom is mostly empty space.
  • The mass and positive charge of an atom is concentrated in a tiny, dense nucleus.

• Planetary Model:
  • Limitations
    1. A nucleus composed entirely of position charges
    2. Could not fully explain the total mass of the atoms (neutrons)

Atomic Models

• Dalton’s Model: Solid, indivisible sphere.
• Thomson’s Plum Pudding Model: Sphere of positive charge with embedded electrons.
• Bohr’s Model: Electrons orbit the nucleus in specific energy levels or shells.
• Modern Quantum Mechanical Model: Electrons exist in probability clouds (orbitals) around the nucleus.

Key Concepts

• Atomic Number (Z): Number of protons in an atom’s nucleus, determines the element.
• Mass Number (A): Total number of protons and neutrons in an atom’s nucleus.
• Electron: Negatively charged subatomic particle.
• Proton: Positively charged subatomic particle.
• Neutron: Neutral subatomic particle.
• Nucleus: Dense, central part of an atom containing protons and neutrons.
• Quantized Energy Levels: The concept that electrons can only occupy specific energy levels within an atom.
• Orbitals: Regions around the nucleus where electrons are most likely to be found, according to the quantum mechanical model.

Additional Points

• Rutherford’s gold foil experiment also indirectly supported the idea that most of the space in an atom is empty.
• While J.J Thomson’s experiment discovered the electron, it was earlier experiments with cathode rays that established the existence of negatively charged particles.

Determining Atomic Properties

Given Protons and Neutrons

If you know the number of protons and neutrons in an atom, you can find:

1. Atomic Number (Z)
   • Equal to the number of protons.
   • Uniquely identifies the element.
   • Use the pe
     Pro
     Con
     Easy way to quantify data

• Bad

Sad

Mad

Not all courses would benifit from this style

riodic table to find the element symbol.

2. Mass Number (A)

   • Sum of protons and neutrons.
   • A = Number of Protons + Number of Neutrons

3. Number of Electrons

   • In a neutral atom, it’s equal to the number of protons (Z).
   • For ions:
     • Positive ion (cation): Number of electrons = Number of protons - Charge
     • Negative ion (anion): Number of electrons = Number of protons + Charge

4. Element Symbol

   • Use the periodic table to find the symbol corresponding to the atomic number (Z).

5. Isotope Notation

   • Represents atoms of the same element with different numbers of neutrons.
   • Notation:
     • A = Mass Number
     • X = Element Symbol
     • Z = Atomic Number

6. Approximate Atomic Mass

   • The mass number (A) is a good approximation.
   • The actual atomic mass is slightly different due to the mass defect and isotopes.
   • Find the precise atomic mass on the periodic table.

Example

• Given:

  • Protons = 6
  • Neutrons = 6

• Deduced:

  • Atomic number (Z) = 6

  • Element Symbol = C (Carbon)

  • Mass number (A) = 6 + 6 = 12

  • Number of electrons (neutral atom) = 6

  • Isotope notation:

      12
      C
       6
    

  • Approximate atomic mass ≈ 12 amu (actual: 12.011 amu)

Key Points

• Protons define the element.
• Neutrons affect the mass number and isotope.
• Electrons determine the atom’s charge.

Isotopes

An atom with the same number of protons but different numbers of neutrons

The mass number is used to differentiate between isotops (e.g. He**-{mass num}** )

N protons = atomic number
Mass number = # protons + # neutrons

Isotopic Abundance

The amount of a given isotope of a n element that exists in nature, expressed as a percentage or decimal.

Calculating

$$ab=\text{percantage in nature}AAM=\%a{b}_1(mas{s}_1)+\%a{b}_2(mas{s}_2)+...\%a{b}_3(mas{s}_n)$$

E.g.

Boron’s atomic mass is 10.81
Isotope Isotopic Abundance
Borem-10: 19.9%
Boron-11: 80.1%

When calculating you need:

• The number of Isotopes
• The mases of each isotope
• The % abundance of each isotope

Calculating relative absundance

• the percent absundance of al isotopes should ad up to 1
• The % abundance of boron-10 is x and the % a abudance of boron 1 is 1-x

---

• Let b-10 = repersent
• let b-1 repersetn 1-x
• $AAM= \x(10.01u) + 1-x(11.0.1u)
• 10.81 = = 10.01x + 11.01-11.01x
• 11.01x - 10.01x = 11.01 - 10.81
• x = 0.20

Ions

Ions
Atoms always want to Get a full shell and become like their closest nobel gas

• Elements that gain electrons:
  • Anion (Non metals)
• Elements that lose electrons:
  • Cation (Metals)
• Elements that share electrons:
  • Molecular

The octet rule

The valence shell must be full as the atom wants to be stable

Polyatomic

An ion with more than one atom
e.g. $P{O}_{4^{3-}}$)

• people tend to think they are dangerous
  • Wrong

Periodic Law

The chemical and physical propeties of elements repeat in a regular, periodic pattern when they are
arranged according to their Atomic Number

Ionization energy

inversely proportionate to ionic radius

• As we go across a period (Left to Right) the atomic radous decreases & EA increases

Taking more electrons away takes increasingly more energy

Electronegativity (EN)

The ability of an atom to attract
the shared Electrons towards itself in a bond

• Cannot be messured experimentally
• Calculated using physical properties
• Values are from 0.7-4

Utils

Bohr Model

Energy Levels

Electron Affinity

• The energy change associated when an electron is added to a neutral, gaseous atom to form an Anion
• Does not apply to Nobel Gasses
• EA Increases from left to right and decreases down a group
• A larger EA value means it’s easier to add an electron to that atom
  • F_{(g)} + e- \rightarrow F^-\,_{(g)} + \text{energy}

Anion

A negatively charged Ion

Cation

A positively charged ion

Defs

Emission Spectrum

Iron’s Emission Spectrum

Nomenclature

• Univalent

All metals in non grouP 1-2

Multivalent:

• Groups 3-12
• Need to include romen numerals

Polyatomic Ions:

Easy way to tell is to find

TO STUDY: OXYACID

Lithium Hydrogen Carbonate

As(C2H3O2)3

Tips and Trics

Acids

Always has Hydrogen in them

• Names don’t need to be exact (e.g. Chlor → Chorine)
  Hypo: 1 oxygen
  Per: 2+ oxygen
  Thio-: Sulfer
• ”Hypo-” is used when there’s one less oxygen than the “-ous” acid.
• 1. Acids formed from polyatomic ions ending in “-ite” are named by replacing “-ite” with “-ous acid”.cx

Binary acid

HX where x is a non-metal

• Hydro-(non-metal)-ic acid

Oxyacids

Look for polyatomic name and add acid

E.G. H2SO4

• Polyatomic: SO4 (sulfate)
• Oxyacid name: Sulfiric
• Final anser Sukfuric acid

Polyatomic name	Oxyacid name	Relative # of oxygens
Hypo…ite	Hypo…ous acid	-2
…ite	…out acid	-1
…ate	…ic acid	reference (before we start gaining/losing)
Per…ate	Per…ic acid	+1

Acid Radicals and Acid Salts

Valance = # of hydrogens being removed
For oxyacids with 2 acidic hyrdrogen (diprotic acids), the bi prefix can be used instead of monohydrogen (e.g. HCO3 can be bicarbonate or monohydrogen)

Hydrates

An ionic compound that have water molecules trapped

• NOT WET
  written with a dot (e.g.)
  $CuS{O}_4\cdot 5H_{2}O$ = Copper (II) Sulfate $\cdot$ Pentahydrate

Molecular Elements

An element

• Always Diatomic

Greek prefixes

1. mono
2. di
3. tri
4. tetra
5. penta
6. hexa
7. hepta
8. octa
9. nona
10. deca

Oxygen drops the first vowel (e.g. monOxide vs monOOxide)

Hydrogen is weird

• NH3 Ammonia
• CH4 Methane

Electronegativity and bondinginherently

To determine if an bond is ionic or covalent you must calculate the difference in the EN values between elements

• Greater the difference ($\Delta EN$): the higher likelihood that the bond will be ionic
  1.7-3.3: Ionic bond
  1.69-0.41: Polar covalent
  Electrons tend to stay in the higher $EN$ atom for longer; unequal sharing
  you have an $⟶$ pointing to the atom with the higher $EN$ value.
  0.4-0: Non-polar covalent
  When two atoms with a $\Delta EN=0$ bond, it’s called a pure bond

Formal Charge

| $\text{Formal Charge}=\text{Normal Valence Electrons}-(\text{Nonbonding Electrons}+\text{Bonding pairs})$

• The difference between valence electrons in an isolated atom and the number of
• For neutral molecules the sum of formal charges must add up to zero
• for ions, the sum of the formal charges must equal the charge of the ion

Quick Scraps

Molecular Polarity

Consider $HCl$ and $H_{2}O$
HCI has 1 polar covalent bond so the molecule has a positive and negative end

How to determine type

1. Determine how many atoms of each element make up the molecule
2. Draw the Lewis structure for the molecule
3. Determine how many covalent bonds there are in the molecule
4. Determine the $\Delta EN$ for each covalent bond in the molecule (polar $0.4>\Delta EN>1.7$ or non polar, $\Delta EN=0.0$ )
5. If there are polar bonds, indicate the partical charges for $\partial -$ for higher $EN$ atom and $\partial +$ for lower $EN$ atom
6. Interpret
   1. Molecule has only one polar covalent bond = polar molecule
   2. Molecule has more than one polar covalent bond, molecule may or may not be polar
      1. Examine the same, is it symmetrical or asymmetrical

NOTE: Polar bonds may not cause the whole molecule to be polar

Polar molecule

• A molecule where uneven distribution of electrons causes one end to be positively charged and the other negatively (POLES!!)

Non-polar molecule

• Has no charged ends

Electronegativity (EN) Increases from LTR. Why?

• The number of protons increases
• Thus, shared electrons attracted more strongly

Electronegativity (EN) Decreases down a group. Why?

• Down a group, the number of protons increase but also more inner electrons, which creates the shielding effect
• Thus, shared electrons are less attracted to the nucleus.
  

Chemical bonding

• Ionic or Molecular

Ionic

• held together by electrostatic forces (negative + positive)
• Arranged in grid like patterns

Patterns

• Metal and non metal
• Hard and brittle
• Relatively high melting and boiling point
  • High due to ionic bonds
• Ability to conduct electricity when dissolved in water (electrolyte)

Properties

• High melting point: Ions are held together by the strong electrostatic forces which are hard to break apart (ionic bonds)
• Hard: The crystal lattice structure is dense and resistant against being stretched
• Brittle: Given that the the Ionic compound is a lattice where $i\pm 1=!i$ on all axis, if you shift the lattice, a positive is next to a positive and a negative is with a negative; this repels ripping the structure apart
• Electrolyte: As the compound dissolves, it releases free flowing ions that can move and carry electric charges

Bonding Capacity

= Max valance - total currently in valence

Formula Units

A repeating structure (in a Ionic compound)

Resonance structure

A structure where you can have a double bond on either side

• Shown via a dotted line or an arrow

Molecular

Two or more non-metals
shared bonds
Poor conductors of electricity

• No transfer of electrons (no ions)
  Low melting point
• No electrostatic forces that need to be overcome

State

• solid
  • Will be a soft solid
  • waxy, flexible or crystalline
• liquid
• gas
  • Most common

Types of synthesis reactions

Two elements will react to form the binary compound

• Metal + Nonmetal >Ionic compound

• Hydrogen + Nonmetal > Simple acid

• Nonmetal + Nonmetal > Molecular compound

Compound + Element > Complex Compound

• Metal chloride + oxygen > metal chlorate

Compound + Compound >Complex Compound

• Two molecular compounds > larger molecular compound

• Metal oxide + carbon dioxide > metal carbonate

• Non-metal oxide + water > oxy acid

• Metal oxide + water > metal hydroxide

Decomposition Reactions

Reverse of Synthesis reactions

Organic

AlkAne are SINGLE BONDS
AlkEne is DOUBLE BOND
AlkYne are TRIPLE BONDS

1. Meth
2. Eth
3. Prop
4. But
5. Pent
6. hex
7. hept
8. oct
9. non
10. dec
11. undec
12. dodec

$C_n{H}_{2n+2}$

Combustion

• Metal combustion
  • Produces basic oxide
• Nonmetal combustion
  • Produces acidic oxide
• Hydrocarbon combustion
  • Complete
  • Incomplete