Unit 1 Chem

These 8 scientists’ research contributed to Atomic Theory and Quantum Mechanics (in order)

1. Empedocles

2. Democritus

3. Alchemists

4. John Dalton

5. J.J Thomson

6. Ernest Rutherford

7. James Chadwick

8. Niels Bohr

Atomic Radius

• The distance from the nucleus of an atom to the outermost energy level.

• As you go across a period from left to right, this decreases – because the number of protons is increasing, and therefore pulling the electrons closer to the nucleus.

• As you move down a group it increases – because additional energy levels are added putting distance between the protons and the outermost electrons. A greater distance means weaker attraction and the atom
  expands.

Ionization Energy

• The minimum energy required to remove the most loosely bound electrons from an atom

• As you go across a period from left to right IE increases - because AR is decreasing electrons are closer to the protons, and electrons are harder to remove.

• As you go down a group IE decreases - because AR is increasing, outermost electrons are further from protons and are easier to remove

Ionic Radius

• A measure of the size of an ion

• Cations are smaller than their neutral atom

  • they lose electrons and therefore a shell, making them smaller

  • The protons on the nucleus gain a greater hold on the electrons because their are fewer of them

• Anions are larger than their neutral atom

  • they gain electrons and complete their outmost shell

  • This means they have more electrons than protons, so the nucleus has a weaker hold on the electrons

  • gained electrons increase the repulsion among electrons

Electron Affinity

• The energy given off of an atom when an electron is gained

• As you go across a period from left to right EA increases, because the AR is getting smaller, so the pull from the nucleus is stronger and nearby electrons are easily attached

• As you go down a group EA decreases, because AR is increasing, and the attraction of nearby electrons to the protons in the nucleus is weaker

Reactivity

• Francium is the most reactive metal - Low IE, lose electrons easily

• Fluorine is the most reactive nonmetal - High EA, gain electrons easily

Electronegativity

• A measure of an atoms ability to attract electrons within a chemical bond

• There is a specific EN for each element

• Trend is the same as ionization energy

• The smaller the AR the the greater the IE, therefore greater EN

• The larger the AR, the lower the IE, therefore lower EN

• unitless scale from 0-4

• Higher EN, greater attraction of shared electrons

Empedocles Theory

• 4 elements: Fire, Air, Water, and Earth

• All substances are made of a combination of these elements

Democritus

• All matter can be divided into smaller and smaller pieces until it is so small it can no longer be divided.

• These tiny particles were called “Atomos“ -meaning indivisible

Alchemists

• Hired by kings to turn base metals, such as copper and lead, into gold

• This was impossible, but they developed glassware and chemical processes that we still use today

John Dalton

• After the Alchemists

• Developed “Dalton’s Atomic/Particle theory“:

  1. All matter is made up of small particles called atoms that can’t be created or destroyed

  2. Each element is made of its own Atom

  3. Atoms of different elements have different properties

  4. Atoms and 2 or more elements can be combined in constant ratios to form new substances

• Dalton’s Model (Solid sphere or bowling ball model)

J.J. Thomson

• After Dalton

• Discovered the Electron by using Cathode Ray tubes

• Plum Pudding model:

  • Atoms are positively charged spheres

  • They have negatively charged electrons embedded throughout

  • Atoms have an overall neutral charge

Ernest Rutherford

• After Thomson

• Nuclear Atom Theory:

  • Gold foil experiment:

    1. Atoms are mostly empty space - proven by most light running through

    2. At the centre of every atom is a small positively charged core - the nucleus - proved by how a few particles bounced back from the gold foil

James Chadwick

• After Rutherford

• Discovered the Neutron:

  • Neutrons have a similar mass to protons

• This explained the different radioactive and mass properties of isotopes

Niels Bohr (State theories, but don’t explain)

• After Chadwick

• Planetary model

• Quantized Energy

• Bright line spectrum

Planetary Model

• Developed by Niels Bohr

  • Electrons orbit the nucleus

  • They exist in specific energy levels of constant energy

  • When electrons gain energy from heat or electricity, they jump from their ground state to a higher energy level (excited state)

  • Electrons immediately fall back to ground state, giving off the the energy as coloured light

Quantized Energy

• Developed by Niels Bohr

  • Electrons are in fixed energy levels or shells found around the nucleus

  • Each level is quantized, meaning they have a specific value of energy

Bright Line Spectrum

• Developed by Niels Bohr

  • A series of bright lines of light produced or emitted by a gas excited by heat or electricity

  • A light spectrum is used as evidence of a new element, since it it is characteristic to each element

Wave - Particle Duality

• This theory explains what keeps atoms from collapsing. (prevents the electrons from being drawn inwards to the protons)

Electromagnetic Radiation

• Travels through space as a wave at the speed of light

• Contains a range of frequencies known as the EM Spectrum

Photons

• Tiny packets of energy that emit EMR

Long wavelength

Low frequency / Low energy

Short wavelength

High frequency / High energy

A quantum of energy

The energy of a photon

Louis de Broglie

• He showed that small particles do not behave like large particles, but contain wave like properties, called the Wave- Particle Duality

Erwin Schrodinger

• This scientist was after Bohr

• He developed wave equations that:

• Describe the energy of the electrons

• Predict the probability of finding electrons in certain regions of space

Arnold Sommerfeld

• He extended Bohr’s work

• Proposed that energy levels are divided into sublevels:

  • S

  • P

  • D

  • F

Werner Heisenberg

• This scientist was after Bohr

• He theorized the probability model/ uncertainty principles

  • Can not predict the the exact location of an electron

  • Can only predict the probability of it being in a certain location

• Electrons act more as waves than particles and don’t lose energy

• There are 3D regions called Orbitals (not to be confused with Bohr’s energy levels or shells)

Orbital

• A 3D space around a nucleus in which Electrons are most likely to be found

• Shape represents electron density. not the path they follow

• Each Orbital can hold up to 2 electrons

Quantum Numbers

• A set of 4 numbers that describe various properties of an orbital

• These numbers are like addresses for locating the position of an electron in an atom

• The 4 numbers are:

  1. Principle Quantum Number

  2. Angular Momentum Quantum Number

  3. Magnetic Quantum Number

  4. Spin Quantum Number

Principal Quantum Number (n)

• This is the first quantum Number

• Uses the symbol “n” to label the energy level

• The larger the “n“, the larger the size of the electron cloud

• 1st level n=1 , 1 sublevel

• 2nd level n=2 , 2 sublevels

• 3rd level n=3, 3 sublevels

• 4th level n=4, 4 sublevels

Angular Momentum Number

• This is the second quantum number

• Uses symbol “I“

• Describes additional electron energy sublevels that form part of the main energy level

• SPDF sublevels each represent a different shape

• S : Spherical - 1 Orbital (2 e-)

• P : Dumbell - 3 Orbitals (6 e-)

• D : Cloverleaf - or elongated dumbbell with a donut around the middle, 5 Orbitals

• F : Complex shapes - 7 Orbitals

Magnetic Quantum Number

• This is the third quantum number

• Tells us the orientation in space of the electron orbit

Spin Quantum Number

• This is the 4th quantum number

• represents the clockwise and counterclockwise spin of electrons

• Electrons occupying the same orbital must spin in opposite directions

• Shown as an up arrow beside a down arrow

Aufbau Principle

This principle states that Electrons fill the lowest energy levels first

Hund’s Rule

This rule states that Electrons remain unpaired as long as possible, meaning the orbitals become half filled before they pair up

Pauli Exclusion Principle

This principle states that there is a maximum of 2 electrons in each orbital, and they must spin in opposite directions

VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory

  • The shape of a molecule is determined by the repulsion between pairs of bonded electrons (like repels like)

  • Non bonding (lone) pairs of electrons around the central atom have a stronger repulsion of bonded electrons

Polar Bond

• Bond contains an EN difference of 0.5 to 1.7

Non Polar Bond

• Bond contains an EN difference of < 0.5

Tetrahedral

• Central atom surrounded by 4 bonding atoms and zero non-bonding pairs of electrons (lone pairs)

• Bond Angle 109.5 degrees

Trigonal Pyramidal

• Central atom surrounded by 3 bonding atoms and 1 non-bonding pair of electrons

• Bond Angle 107 degrees

Trigonal Planar

• Central atom surrounded by 3 bonding atoms and zero non-bonding pairs of electrons

• Bond Angle 120 degrees

Bent

• Central atom surrounded by 2 bonding atoms and 2 OR 1 non bonding pairs of electrons

• 104.5 degrees

Linear

• Central atoms surrounded by 2 bonding atoms and zero non bonding pairs of electrons, OR no central atom

• Bond Angle 180 degrees

Trigonal Bipyramidal

• Central atom surrounded by 5 bonding atoms and zero non-bonding pairs of electrons

• Applies to Phosphorus when it bonds to 5 atoms (Octet rule exception)

• Bond Angle is 90 and 120 degrees

Octahedral

• Central atom surrounded by 6 bonding atoms and zero non bonding pairs of electrons

• Applies to Sulphur when it bonds to 6 atoms (Octet rule exception)

• Bond Angle 90 degrees

Octet Rule Exceptions

• Hydrogen: 2 ve-

• Boron: 6 ve-

• Beryllium: 4 ve-

• Phosphorus: 5 bonds = 10 ve-

• Sulphur: 6 bonds = 12 ve-

Metallic bonds

• This bond is formed from electrons being removed from positively charged ions (metals)

• The ions pack closely together, while their valence electrons behave like a mobile sea of electrons

• The electrons move throughout the metal

• the electrons are called delocalized electrons and they don’t belong to any specific atom

• The electrons act as a glue

• More electrons = stronger “glue“, higher melting point and hardness

Alloys

• Mixture of metals with other metals or sometimes non metals

• Alloys are often stronger and harder than the original metal

• Added metal may provide extra valence electrons

• Different sized alloy metal atoms make it more difficult for metal layers to slip off one another

• Two types:

  • Substitutional

  • Interstitial

Substitutional Alloys

• A type of Alloy

• Atoms of added metal can be substituted for those of another if they are approximately the same size

• Ex. Brass= Copper + Zinc

Interstitial Alloys

• Atoms of added metal are small enough to fit in the space between the other metal atoms

• Ex. Steel = Iron (Black) + Carbon (Red)

The 2 groups of Chemical bonding

1. Crystal Lattices

2. Small Molecule Interactions

Crystal Lattices and their 3 types

This type of chemical bond has a very large number of atoms bonded together in a regular, repeated sequence, usually 3D.

The three types are:

1. Ionic

2. Covalent (Network Solids)

3. Metallic

Ionic Crystals

• This crystal is formed from ionic bonds

• There is an attraction between +ve and -ve in the lattice

Metallic Crystals

• This crystal is formed from metallic bonds

• +ve ions held together by delocalized electrons in the lattice

Covalent Crystals

• This crystal is also called a network solid

• they are formed from covalent bonds

• atoms held together by pairs of shared electrons in lattice

Small Molecules and its two types of forces

This type of Chemical bond contains forces that attract one small molecule to another.

The two types of forces are:

1. Intramolecular

2. Intermolecular

Intermolecular Forces and its types

• This type of force creates a weaker force of attraction between separate molecules

• Weaker than Intramolecular forces

• They hold adjacent Covalent molecules together

• Exist especially during liquid or solid state

• The three types are known as Van der Waal’s Forces:

  1. London Dispersion

  2. Dipole - Dipole

  3. Hydrogen Bonding

London Dispersion force

• A Van Der Waal force

• The constant vibration of electrons creates temporary dipoles in the molecule

• The force of the temp dipoles increases as the molecule increases in size

• Caused by a shift in the environment

• Found in small molecules

• If the charge is produced in one molecule, electrons can be attached

• LD is short lived but frequent

• Only effective over short distances

• Relatively weak bonds

• Strength is affected by larger molecules, which have more electrons and more temporary dipoles (becomes stronger)

Dipole - Dipole force

• Found in polar molecules with a small permanent dipole

• The negative end of one molecule is attracted to the positive end of another molecule

• Relatively weak bonds, but stronger than LD

• Strength is affected by EN difference in bonding atoms

Hydrogen Bonding

• A strong dipole- dipole force between H and O, N, or F of another molecule (High EN difference)

• Found in polar molecules with large permanent dipole (F, N, O)

• Relatively strong bond

• Strength is affected by EN difference in bonding atoms

• Strength affected by number of H-N, H-F, and H-O bonds in each molecule

Rank the lattice structures and Van der Waal’s forces from highest to lowest bond strength

1. Covalent (Intramolecular)

2. Ionic

3. Hydrogen (Intermolecular)

4. Dipole - Dipole (Inter)

5. London Dispersion (Inter)

Vapour Pressure

The rate at which a liquid evaporates

Stronger forces =

lower rate of evaporation and vapour pressure =

Low boiling point =

High vapour pressure =

Higher boiling point =

Low vapour pressure =

Melting and Boiling Point

• These points are measures of the strengths of the bonds/intermolecular forces

• It involves the breakdown of the lattice or organized structure between covalent molecules.

Allotropes

• Different physical forms of the same element

• Ex Carbon

  • Diamond

    • 3D

    • Consists of many carbon atoms, such that each central carbon atom is surrounded by 4 others in a tetrahedral