Chemistry Unit One exam

yay!

Spectroscopy - Higher energy levels means

Electrons have absorbed energy so they are excited to a higher energy level.

Spectroscopy - Higher energy levels are unstable therefore.

They return to ground state re-emitting the excess energy, in some elements it will be re-emitted

Quantum mechanics - types of sub shells

S P D F

Quantum mechanics - Subshells holding

Subshells can only hold 2 electrons, when p has 6 electrons in its subshells, in reality it has 3 orbitals that can be filled

Quantum Mechanics - Shapes of orbitals

no definite shape however there are places where electrons frequent

S = Sphere

P = Dumbbell

D = Clover

F = Complex

Quantum Mechanics - Shell order

All 3 shells need to be together all 4 shells need to be together and so on. For example: 1s²,2s²,2p^6, 3s², 3p^6, 3d^10, 4s², 4p^5

Quantum Mechanics - Special rule of filling

Hypothetically, if you had Copper with 4s² and 3d^9 this is unstable so we move one from the 4s² to the 3d^9 in order to gain more stability within the orbitals. Better to have half filled shells than somewhat filled shells.

Conversion

bigger to smaller multiply - smaller to bigger divide

The Periodic Table - Mendeleev’s

Ordered by valency and by properties.

Quantum Mechanics - Nobel gases

They cannot have incomplete shells

The Periodic Table - The modern periodic table

arranged by atomic number and similar properties (Like Mendeleves’s version)

Trends In The Periodic Table - Atomic Radius

Distance between from the centre of the nucleus to the outermost shell.

Going down in the periodic table increases atomic radius because the number of shells increase. Additionally, distance between nucleus + valence electrons = lower force of attraction. Weaker electrostatic attraction.

Going across a period the atomic number increase and the number of shells stay the same so there is stronger electrostatic attraction, therefore there is a decrease in atomic radius because of the pull between the nucleus and the electrons

electrostatic attraction = the pull between the nucleus and the electrons (opposites attract)

Trends In The Periodic Table - Core Charge

The pull on the valance electrons from the nucleus (related to the atomic radius). Core charge increases across a period and is a constant within a group. Decreases as the distance between valance electrons and nucleus decreases the core charge.

Core charge = charge on the nucleus - the charge of the inner shell electrons

Trends In The Periodic Table - Electronegativity

How well a nucleus attracts other electrons (bonding). The higher the attraction between the nucleus and valence electrons the greater the electronegativity.

Electronegativity relates to -

core charge: as the nuclear charge increases across a period and the electron shielding remains constant so the electrostatic force increases therefore electrons are more attracted. Electronegativity increases across a period.

atomic radius: A distance between the nucleus and the outer-shell, as the distance increases the attractive forces decrease and electrons are held weaker. Therefore electronegativity decreases down a group

Trends In The Periodic Table - First Ionisation

Ionisation energy is the energy required to remove an electron from an uncharged atom. The higher the attraction means more energy needed to remove it.

Group 1 has low ionisation levels as they can lose one electron to gain a stable out shell easily.

Noble gases have very high ionisation as they already have a stable octet

This relates to electronegativity. If there is a stronger pull then its harder to remove an electron

Trends In The Periodic Table - First Ionisation Energy with Core Charge and Radius

Low core charge + small radius = electron held fairly strongly

Low core charge + large radius = electron held weakly

High core charge + small radius = electron held extremely strongly

High core charge + large radius = electrons held neither strong or weakly

Trends In The Periodic Table - Electronegativity Difference in Terms of Polar and Non-Polar

0 = non-polar covalent

0=>0.4 = slightly polar covalent

0.4=>1.7 = polar covalent

greater than 1.7 = ionic

Trends In The Periodic Table - Metallic Characteristics

Lovely Heavy Ducks Make Daily Morning Beeps

lovely = lustrous

heavy = hard

ducks = ductile

make = malleable

daily = dense

morning = melting point (high)

beeps = boiling point (high)

Trends In The Periodic Table - Metalloid Characteristics

A B C

A = always a semiconductor

B = Brittle

C = Conduct heat and electricity

Trends In The Periodic Table - Non-Metallic Characteristics

PIGEON

P = Poor conductors

I = Ionisation high

G = Gain electrons

E = Electronegativity high

O = Oxidising

N = Non malleable, non ductile, non lustrous

Trends In The Periodic Table - Reactivity

From the outside in it goes from reactive to least reactive because of electronegativity the highest and lowest will always be reactive. Noble gases are inert so they do not react

Covalent Bonding Non Metals - Forms

2 forms: gase and liquids

Covalent Bonding Non Metals - Molecular bonding

Sharing of valence electrons so that each atom has a stable outer shell

Covalent Bonding Non Metals - Properties of molecular substances

Low melting and boiling point and don’t conduct electricity. This is due to weak forces of attraction and no freely moving charged particles. Atoms that share electrons have high electronegativities so they attract atoms easily. It requires a lot of energy to separate covalent bonds due to having a stable octet.

Covalent Bonding Non Metals - Intramolecular bonding

within the molecule (eg. covalent bonding, metallic bonds and ionic bonding). Can only be broken by a chemical reaction

Covalent Bonding Non Metals - Bond length in terms of attraction

the stronger the attraction the shorter the bond length.

Energy: single bond < double bond < triple bond

Bond Length: single bond > double bond > triple bond>

Structures - Different forms

Electron dot diagram/lewis structure: includes lone pairs in the form of dots (DOTS ONLY)

valence structure: shows all bonds including lone pairs with lines (LINES ONLY)

Structural formula: shows bonds between atoms as lines

Space filling: 3D display showing it as one compound together

Ball and stick: 3D structure showing atoms and bonds as lines

Covalent Bonding Non Metals - Naming Covalent Compounds

1. First element named in full

2. second element in the formula as if it was anion. The name is shortened and given an “ide“ suffix

3. A prefix is used to indicate the number of each type of atom in the molecule

4. elements names beginning with a vowel may have some letters dropped if the prefix ends with a vowel

eg. Carbon monoxide

dinitrogen pentoxide

nitrogen tribromide

Covalent Bonding Non Metal - Polar and Non-Polar Covalent Bonds

Non-polar: symmetrical or equal electronegativity

Slightly polar covalent: 0-0.4 difference in electronegativity

Polar bonds: asymmetrical and 0.4-1.7 difference in electronegativity

Anything greater than 1.7 is ionic

Covalent Bonding Non Metal - Forces Between Molecules/Intermolecular forces

Between molecules (dispersion forces, dipole-dipoles, hydrogen bonds) these can be separated by heating the substance

Covalent Bonding Non-Metals - Dispersion Forces

The electrons within an atom or molecule are constantly moving. This can create a slightly negative charge as it rotates this also allows a side to be slightly positive as well. Creating an instantaneous dipole. Dispersion forces are non-directional. The larger an atom or molecule the more electrons hence more dispersion forces (very weak).

Covalent Bonding Non Metals - Dipole Dipole Attraction

Polar molecules have permanent dipoles. Dipole-Dipole attraction is non-directional as the molecule can move around. Weak intermolecular forces. The more polar molecules are the stronger the dipole-dipole attraction.

Covalent Bonding Non Metals - Hydrogen Bonding

When some of the most electronegative elements bond with hydrogen (FONCL). Since the more electronegatively charged atoms pull the electrons it makes the hydrogen more positive. Hydrogen bonding is the strongest out of the intermolecular bonds. It has higher melting and boiling points due to its strong attraction.

Covalent Bonding Non Metals - Carbon structure and bonding

there are three main allotropes (forms of pure carbon): diamond, graphite and amorphous carbon (without structure). There is network lattice structures, layered lattice structure and the no structure

Reaction of Metals - Bonding Model Explanation High Melting and Boiling Points

Strong electrostatic forces of attraction between positive cations and the sea of delocalised electrons holds the metallic lattice together

Reaction of Metals - Bonding Model Explanation Metals are good conductors of electricity

In an electric circuit, free moving delocalised electrons will move towards positive electrodes (conductive materials) and away from the negative electrodes

Reaction of Metals - Bonding Model Explanation malleability and ductilibility

When a force causes layers of metal ions to move past each other, layers of ions are still heald together by the delocalised electrons between them

Reaction of Metals - Bonding Model Explanation Good conductors of heat

when delocalised electrons bump into each other and the metal ions, they transfer energy to each other. Heating gives the ions and electrons more rapidly. The electrons are free to move and os they transmit this energy rapidly through the metal lattice.

Reaction of Metals - Bonding Model Explanation Lustrous

The free moving electrons cause the metal lattice to reflect light and appear shiny

Reaction of Metals - Bonding Model Explanation Densities

The ions in a metal lattice are closely packed. The density of a metal depends on the mass of the ions, their radius and their pacjjing arrangement in the lattice

Reaction of Metals - Limitations of the Metallic Bonding

Doesn’t adequately explain:

• differences in electrical conductivities

• Wide range of melting and boiling points

• magnetic properties of some metals

The model cannot represent:

• Movement of electrons

• Vibrational and rotational movement of particles

• Scale of particles

• atoms are mostly empty space not solid balls

Reaction of Metals - Reactivity

Reactivity of metals increases down a group and decreases across a period. Reactivity of metals is related to how readily they lose electrons to form positive ions (Ionisation)

Linear economy

In a linear economy substances are produced, used and discarded

Circular Economy

In a circular economy substances are used, recovered and recycled to use again

Ionic Bonding - What is the bonding between

Metals and non metals. Charged particles that form when an atom gains or loses one or more electrons. Cations and Anions bond together and transfer electrons to each other to gain a stable octet. Ionic bonding is directional as it only happens in a direct line between the adjacent oppositely charged ion

Difference between Covalent and Ionic

covalent share the electron while ionic give the electron

Ionic Bonding - Properties

• high melting temperatures (because of large amounts of energy since they become stable)

• forms hard and brittle crystals

• solids do not conduct electricity

• molten does conduct electricity

• when dissolved in water it does conduct electricity (because it allows the charged particles to move freely)

Different types of bonding

Metallic - two metals

covalent - two non-metals

ionic - metal and non-metal

Chromatography - Dissolving

Like dissolves like

Chromatography - Polarity and chromatography

Variation in polarity mean that components will be more or less strongly attracted to different solvents

Chromatography - Rf Values

Rf = distance moved by components/distance moved by solvent(total distance moved)

Quantifying Atoms and Compounds - Relative Atomic Mass (RAM)

RAM is a weighted average of all the naturally occurring isotopes of an element.

RAM = (RIM1 x %) + (RIM2 x %) + …./100

(RIM is relative isotopic mass)

RAM is also known as Ar

Quantifying Atoms and Compounds - Relative Molecular Mass (RMM)

Symbols: Mr or RMM

Calculated by using the mass of atoms eg. Water = 2× 1.0 × 16.0 = 18.0

Quantifying Atoms and Compounds - The mole

Avogadro’s number = 6.02 × 10²³

Quantifying Atoms and Compounds - How to find the Mole

n = N/Na

n = number of moles

N = number of particles

Na = Avogadro’s number

n = m/M

n = number of moles

m= mass of substance given (g)

M = Molar mass (gmol-1)

Quantifying Atoms and Compounds - Molar mass (Mr)

The mass of one mole of the substance.

Finding M => NH3 = Mr(N)= 14.0 Mr(H)= 1.0

Mr(NH3)= 14.0 + 3 ×1.0 = 17.0

Quantifying Atoms and Compounds - Theoretical Percentage Composition

Percentage of the mass of a compound that is contributed by each element.

% of one element = total molar mass of atoms of the element in one molecule/Molar mass of the molecule x 100

eg. H20 H= 2.0/18.0 × 100

Quantifying Atoms and Compounds - Empirical Formulae

Putting molecular formulas to their simplest forms

m(atom)/M(atom) : m(atom)/M(atom)

divide both sides by the smaller number

Round to the nearest whole number

Quantifying Atoms and Compounds - Hydrated Salts

A hydrated salt is an ionic compound which contains water molecules in solid crystal the degree of hydration

Organic Chemistry - Types of Formulas

Structural - diagram of the bonds between atoms

Condensed(semi structural formula) - eg. CH3CH2CH3

Molecular formula - eg. CH4

Skeletal - shows everything as lines

Organic Chemistry - Properties of hydrocarbons

• Intramolecular bonding: non-polar covalent

• Intermolecular bonding: dispersion forces

• non polar

• MP and BP increase (carbon chain increases in length therefore there is an increase in strength in dispersion forces)

• Branching lowers MP and BP (because molecules cannot pack closely together so dispersion forces are weaker)

• Hydrocarbons are insoluble in water (because they are polar)