Hitting the Jackpot

Atoms

Subatomic particles

Atomic Number: Number of protons in the nucleus

Mass number: The number of nucleons

Dalton’s Atomic Theory

Isotope

Chemical properties-depend on the number of electrons, since elements have the same number of electrons, their chemical properties are identical

Physical properties: Depends on the neutrons of an atom, so the physical properties of isotopes pairs are slightly different’

Natural Abundance-Refers to the proportions of isotopes of a chemical elements as naturally found on a planet

Atomic mass is the relative percentage of all the naturally occuring isotopes of that element

Mass Spectra

How a mass spectrum for an element or molecule is created

1. A sample put into the spectrometer and vaporised;

2. An electron beam creates ions with a 1+ charge

3. Ions are accelerated by an electric field and focused

4. Ions are deflected by an electromagnet. Changing the strength of the electromagnetic changes the deflecting pattern based on ion mass

5. The ions are detected by the detector and the mass spectrum plotted

Emission spectra

Light: A form of electron magnetic radiations, a wave of it described with frequency and wavelength

Spectrum

A spectrum is produced when a light source (sunbeam, torch, laser etc) passes through a transparent prism (piece of glass, or a diffraction grating) and the light is dispersed through an angle that depends on the wavelength of the light passing through.

Continuous spectrum: 

Contains all possible wavelengths (white light) then a continuous spectrum results

Line Spectra:

Only contains a few light colors surrounded by black. At higher frequencies they converge.

Emission spectrum of hydrogen

When electrons absorbs, heat or energy the go up energy levels, this is unstable so after some times they go down energy levels and emit photons at a specific wave length. Higher energy levels are closer together on the graph as they are closer on together in the atom, thus requiring less energy to move up and down. Each energy difference then refers to a specific wavelength. Use to prove the existence of energy levels

The energy of an electron in an energy level is equal to:

The energy emitted (or absorbed during energy level transition is equal to:

Quantum-Mechanical Model

Heisenberg uncertainty principle showed us that we cannot know both location and the velocity of very small particles such as the electron

Energy Levels and Orbitals

7 Energy levels, 7 periods of periodic table

S block:

1. First two groups of the periodic table

2. Spherical and only one orientation

3. So can only hold two electrons

P Block:

1. last 6 groups of periodic table (13-18)

2. Can hold 6 electrons as it can have three different directions (x, y and z)

D Block:

1. Transition metals

2. Contains 5 sub levels so maximum of 10 electrons

F Block:

1. Contains 7 atomic orbitals

2. Lanthanide and actinides

First energy levels can hold 2 electrons, then 8, 18, 32

Transition metals d block

Group 13-18 p block

Notable principles:

1. The Aufbau principle states that when adding electrons to an atom, the lower energy orbitals must be filled first. (s, p, d, f)

2. The Pauli exclusion principle states that an atomic orbital can only hold two electrons and they must have opposite spins.

3. Hund’s rule states that when we have degenerate orbitals (orbitals of the same energy) then each orbital is filled with a single electron before being doubly occupied turning them into a

4s sub level fills before 3d as it is lower in energy. However, 4s level electrons are taken away before 3d because 3d has lower energy.

Exceptions:

Every element in the period of Chromium and Copper have one less 4s electron but one more 3d electron. But then when leaving

Cu: 1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰

Cr [Ar] 4s¹ 3d⁵

Electron Configuration

Electron configuration is determined by distributing the atom’s electrons among the levels and sublevels based on a set of stated principles and a rule:

Aufbau Principle

An electron occupies the lowest energy orbital that can receive it. from s, p, d, f. This means that first 4s is filled as it is lower than 3d

Pauli Exclusion Principle

No two electrons in the same atom can have the same set of four quantum numbers.

Hund’s Rule

Orbitals of the same energy (degenerate orbitals) are occupied by one electron before any orbital is occupied by a second electron, and all electrons in separate singly occupied orbitals must have the same spin.

Exceptional Electron Configuration:

Chromium and copper have exceptions. Instead of having 4s² one of the electrons is giving to d block so it can have half filled and fully filled blocked respectively

How to do it:

1 Fill up the electrons from the lowest energy level.

2 Put the electrons in each sub-energy level one at a time with spin up then pair up any more electrons added with spin down.

First IE Energy

Energy ionization energy is the energy required to remove one mole of the most loosely held electron from one mole gaseous atoms to produce to produce one mole gaseous ions each with an oxidation state +1

Equation for Ionisation energy:

M(g) → M⁺(g) + e^–

IE is governed by:

• the charge on the nucleus

• the amount of screening by the inner electrons

• the distance between the outer electrons and the nucleus

As you go down a group the increase in nuclear charge is the same as the increase in the number of inner electrons. However as you go down distance between nucleus and outer electrons increases so first ionization energy decreases.

IE goes up and down as removing an electron from a half filled or fully filled block is harder. But, overall there is an increase in IE as more neutrons creates stronger attractions pulling in electrons. 

Electrons at higher energies are easier to remove as they are farther away from the nucleus

Ion formation

Ionic bonding

Ionic Compounds

Ionic compounds properties:

• High melting and boiling points, due to high electrostatic attraction and crystal structure

• Hard and Brittle as solids, due to hard electrostatic attraction however brittle as when hit positive and positive ions are repulsed overcoming electrostatic atraction

• Non conductive when solid but conductive when liquid (molten or aqeous)

• When dissolved in water called hydration

Lattice Enthalpy:

Definition: The energy required to convert one mole of solid ionic compound into gaseous ions.

Stability: Depends on all of the attraction between the actions and anions

The more energy released during the formation of the (solid), the more stable the solid is.

Strength depends on:

1. The charge on the ions, higher charge means, greater attraction between them and strong lattice enthalpy. More important than size

2. Size of the ions, smaller the atoms the stronger attraction between them and the stronger lattice energy

Metallic Bonding

A metal contains positive metal ions surrounded by a sea of delocalized electrons.

The smaller the cation and the more electrons in the ‘sea’ the stronger the metallic bonding and the higher the melting point

As you go down the group bonding strength decreases as atomic radii increases thus lower melting point

Across a period nuclear charges increases and atomic radii is constant so metallic bonding becomes strong so boiling point increases

Transition metals good at conducting electricity due to delocalized electrons

Malleability:

1. Be able to be compressed into shape without being broken

Ductility:

1. being able to be strecthced out

Transition elements

The periodic table

Periodicity

The atomic radius is half the distance between two nuclei of a diatomic molecule

Influenced by three factors:

1. Energy Level-Higher energy level is further away from nucleus

2. Charge on nucleus, more charge pulls electrons in closer

3. Shielding effect In a group, you add electrons for each level, this makes teh radii larger as you go down the table

Across a period

1. Number of protons increases so increase in nuclear charge

2. Electrons are filled in the same energy levels

3. So attraction increases thus atomic radii decreases

Down a group:

1. Number of energy levels increase

2. Protons increases, however the increase nuclear charge is exactly offset by electron shielding

3. Thus atomic radii increases

Ionic Radii:

The ionic radius is the radius of the most common ion formed

Positive ions:

When a cation is formed, an electron is removed, decreasing repulsion between electrons. Thus can be closer together, thus smaller radius

Negative ions:

when an anion is formed, an electron is gained. This increases the repulsion between electrons. The larger repulsion means a larger radius

Electronegativity:

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons

As the atoms get larger, any bonding pairs gets further and further away form the nucleus and so is less strongly attracted towards it. In other words, as you go down the group the elements become less electronegative.

Electron affinity:

Energy relased when an electron, in addition to the existing electrons, is attached to a neutral gaseous atom or. molecule. Similar trend to that it increases form left to right. Electron affinity decreases as you go down group 1 but the trend is not the same for all groups. If addition of an electron creates a less stable electron configuration the energy relased EA is less

Alloys

The periodic table

Periodic Law:

When elements are arranged in order of increasing atomic number their physical and chemical properties show a identifiable pattern

Group:

Also know as columns. Elements have similar characteristics. Is a series of elements that have the same valence electron number.

Period

The chemical properties gradually change from reactive metals ot reactive non metals

it is a sequence foofelements in which the same energy level is being filled

Is the number of energy levels that contains electrons.

Group 1 and 17 elements

Alkali Metals:

• Reactivity increases down a group because the distance from teh nucleus to the valence electrons increases as you go down the group, the energy needed to remove that electron gets lower. This results in the reactiity increases as you go down a group.

Halogens:

• Always diatomic molecules

• Going down the group electron affinity decreases thus become more non metals and less reactive

Melting point of group 17:

As it goes down the group melting point increase because mass increase equals increase number of proton and electrons equals increase of IMF.

Oxides

Oxidation: Is when a species loses electrons

Reduction: Is when a species gains electrons

Acid Rain

Acid deposition refers to the acidic particles and liquids that deposit or fall to the Earth. It includes liquid deposition of acidic gases such as NO_x and SO₂ by rain, hail and snow solid deposition fo acidic gases and particles.

Rain is naturally acidic with a pH around 6. This is because of the high concentration of CO₂ in the atmosphere.

Rule for determining oxidation state

Charge vs Oxidation state: Charge is the loss or gain of electrons, oxidation state that number are assigned due to a set of rules allowing us to know if oxidation or reduction has happened.