CHM1011 Full Notes PDF

Preview

Summary

Complete notes of entire semesters content from workshops and pre workshop readings...

Description

Week 1: Introduction to Chemistry

Weekly objectives • • • • • • • •

Understand the structure of the atom and important discoveries that revealed the existence of protons, electrons and neutrons. Calculate the mole, molar mass, mass, concentration and dilution Balance chemical equations and determine the empirical formula, percentage yield and limiting reagents Recognise ionic bonding and name common salts Recognise covalent bonding and bond length Understand and apply the Aufbau principle, Hund’s rule and Pauli exclusion principle Be able to write out full and condensed electronic configurations Identify and explain the origin of special cases in electronic configurations

Element: matter containing atoms all with the same number of protons. Molecule: matter containing atoms bonded together in a definite structure. Molecular formula: a formula that shows the number of each type of atom in a molecule. Compound: matter containing two or more elements in definite proportions. Mixture: matter containing more than two compounds and/or elements.

The types of bonds There are two types of bonding you will need to know for this semester: -

ionic covalent

Ionic bonds Ionic bonding is the electrostatic attraction between a cation (positively charged atom) and an anion (negatively charged atom). The ionic bond forms between the cation and the anion - forming a salt. The overall charge of the salt needs to be zero. Salts tend to have high melting points and can have many colours

Covalent bonds -

-

-

Covalent bonds are where atoms share their electrons to fill their outer shell. Elements which partake in covalent bonding are non-metals (B to F, Si to Cl, As to Br and I) as well as hydrogen. When two atoms join together to form a covalent bond, you need to consider all the electrostatic interactions. There is attraction between the opposite charges (negatively charged electrons and positively charged nuclei) and repulsion between the electrons, and between the nuclei The most stable arrangement of atoms occurs when the potential energy is an absolute minimum. We view the electrons as being shared between the nuclei and call this shared electron density a covalent bond

-

Homonuclear molecules have all the same element covalently bonded together. e.g. F2, H2, S8, C60 Heteronuclear molecules have more than one element covalently bonded together. e.g. CH4, N2O, C7H8 Bond length is the separation distance at which the molecule has the minimum energy over the separated atoms.

Chemical reactions A chemical reaction is the production of completely new substances from one or more molecules. We can observe chemical reactions by: -

changes in colour or temperature evolution of heat changes between the states (solids/liquids/gases) emission of light appearance or texture of a solid substance

A phase change is not a chemical reaction. Common chemical reactions you will be studying this semester include: -

Combustion Acid - Base Precipitation - dissociation Redox: Oxidation/reduction

Stoichiometry -

-

Avogadro’s constant= 6.022 × 1023 One mole contains exactly 6.022 × 1023 elementary entities Molar mass (grams/mol) M=m/n

Concentration

-

The definition of concentration is the amount of solute dissolved of a particular volume of solution. n = cV

Dilution - The number of moles of the solute remains the same when we dilute the solution. c1V1 = c2V2

Orbitals Ground state – the lowest possible energy state for a set of electrons Excited state – when one electron is at a higher energy level to its ground state. Electrons require energy to move to an excited state Degenerate – a set of orbitals which have the same energy Aufbau principle -

'The Aufbau principle' literally translates to 'the building-up principle’, and that’s exactly what we do - placing each electron in the lowest energy, permittable level.

Hund’s rule -

Degenerate orbitals in the ground state: • electrons cannot be paired until each orbital in the set contains one electron • single electrons must have parallel spins

Pauli principle -

Each electron has their own unique orbital “name” (quantum numbers) No two electrons have same set of quantum numbers

Electronic configurations of cations and anions Cations

- Determine the electronic configuration of the neutral element - Remove the valence electrons to equal the charge on the cation Anions

- Determine the electronic configuration of the neutral element - Add the extra electrons to equal the charge on the anion

Week 2: Understanding the Periodic Table

Weekly objectives • • • • • • •

Discuss the reasons for periodicity with respect to the periodic table of elements. Define the properties of ionisation energy, electron affinity, atomic radius. Identify the trends in the periodic table of elements for ionisation energy, electron af Identify patterns in successive ionisation energy and appreciate the implications with configurations Predict the size of anions and cations in relation to their neutral atom Understand the concept of electronegativity and the trends in the period table. Understand the implications electronegativity has on covalent and ionic bonding

Periodic trends Zeff (effective nuclear charge) -

net charge of the nucleus after the attraction and repulsion of all the electrons are taken into account. The effective nuclear charge is the difference between the attraction and repulsion. There are two general rules to remember about Zeff: • as you work across a row of the periodic table, the number of protons increase but the number of shells remain the same. Thus Zeff increases from LHS to RHS across the periodic table. • as you work down a column of the periodic table, both the number of protons increase and the number of shells too. The result is a decrease in Zeff from top to bottom of the periodic table.

Atomic Radii -

The number of shells; within a group, the number of shells is the determining factor • atomic radius generally increases in a group from top to bottom. Zeff; within a row, Zeff is the determining factor • atomic radius generally decreases in a period from left to right

Ionisation energy -

Ionisation energy is the amount of energy required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ions •

• • -

-

As atomic radii decreases, it take more energy to remove an electron, since the electrons are closer to the positively charged nucleus. This electrostatic attraction requires more energy to overcome it. Ionisation energy generally decreases down a group as the outer electrons are furthest from the nucleus Ionisation generally increases across a period as Zeff increases.

In summary, ionisation energy trends correlate with atomic radii trends. Thus the larger the ionisation energy (IE) to remove the first electron, the smaller the atomic radius. Atoms with lower values of ionisation energy tend to form cations during reactions, whereas those with a high ionisation energies (except noble gases) often form anions.

Successive ionisation energies One the first electron has been removed from an atom, other electrons can then be removed. These are called 2nd ionisation (IE2), 3rd ionisation (IE3), 4th ionisation (IE4) and so forth. The energy required to remove each of these subsequent electrons increases dramatically as the ion becomes more positive and the attraction of valence electrons to the nucleus is larger. There is a drastic jump in energy when removing a core electron because atomic radii is much smaller and harder to remove an electron. For example: IE2 = second ionisation energy: removes a second electron from the gaseous cation: ion+(g) → ion+2(g) + e- IE2 > IE1

Electron affinity

-

The energy change accompanying the addition of 1 mol of electrons to 1 mole of gaseous atoms or ions Atom (g) + e -→ ion- (g) Electron Affinity (EA1)

-

can be negative or positive values. ion- (g) + e - → ion2- (g)

-

Subsequent Electron Affinity values are always positive. Electron affinity increases across a period, yet there is no apparent trend down a group

Ion size Cations

-

Radii for atomic cations always smaller than neutral atom

Anions

-

Radii for atomic anions are always larger than neutral atom

-

The relative ability of a bonded atom to attract the shared electrons

-

•

as you move down a group on the periodic table, the electronegativity of an element decreases because the increased number of energy levels puts the outer electrons very far away from the pull of the nucleus.

•

Electronegativity increases as you move from left to right across a period on the periodic table

Difference between electronegativity of two elements determines whether bonds are mostly covalent or mostly ionic

Week 3: Orbitals, orbitals everywhere... Weekly objectives • • •

• • • • • • •

Recognise light as one form of electromagnetic radiation.

Understand the meaning of the terms frequency and wavelength, and their interrelationship νλ = c. Investigate the concept of “wave-particle duality”, as revealed through the profound ideas and experiments of Planck, Einstein, Bohr, DeBroglie, Heisenberg and Schrodinger - all Nobel prize winners! To see that light has both wave and particle-like properties (hence photons) and that the same is true for electrons. To understand the need for a new “quantum mechanics” to describe the behaviour of electrons in an atom; this replaced the “classical mechanics” used since the time of Newton. Discuss the use of the four quantum numbers Relate quantum numbers to electronic configurations and trends in the periodic table Explore the concepts of penetration and shielding & explain the origin of nondegeneracy between the orbitals of many electron atoms. Define an atomic orbital & discuss the different shapes and features Define nodes and recognise the nodes for s-, p- and d-orbitals

Electromagnetic radiation -

-

All radiation travels at the same velocity. This velocity is called the speed of light (c), and is 3.00×108 m s-1 Radiation can be measured as: • wavelength: symbol- λ and units- metres • frequency: symbol- ν and units - Hz, or s-1 The relationship relating frequency and wavelength is νλ = c

-

Amplitude, or intensity of the radiation, is the measure of the strength of the light - eg the power on laser beams.

Relationship between wavelength, frequency and energy -

Energy comes in packets called quanta. Quanta, the plural of quantum, is basically the minimum amount certain properties of a system can change In relation to light, the quanta is known as photons, particle-like interpretation of light waves. Einstein used Planck’s idea to concluded that photons have energy proportional to frequency: E = hν

where h is Planck’s constant, 6.63×10-34 J s. Discovery of Quanta using the Photoelectric effect -

-

The experimental setup that led to the Light was directed onto a metal plate Voltage was measured to determine if metal Only after exceeding a certain measured voltage This ejection of electrons from the amount of energy (aka frequency) Proved that light had properties of

Absorption is an electron absorbing energy to go to a higher energy level. Emission is an electron losing energy, releasing that energy as emitted light Gas absorbing light at specific wavelengths, causing the dark lines on the spectrum. Gas emitting light at specific wavelengths, causing the coloured lines on the spectrum.

Atomic emission of light -

Electrons in an atom can only occupy certain orbits corresponding to certain energies. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another.

Wave particle duality -

Louis de Broglie proposed that if light can have material properties, matter should exhibit wave-like properties. de Broglie stated that a moving particle has a wavelength given by λ = h/p, where p is linear momentum based on mass and velocity of the particle. • •

-

Fast electron, 5.9x106 m s-1, λ =1.2x10-10 m Usain Bolt, 10 ms-1, λ =8.3x10-37 m !!

Main point: • For very light things (eg. electrons), λ is significant! • But for macroscopic objects λ is far too small to be detected.

Matter waves and the electron So it was proposed that electrons have wavelike motion and are restricted to orbits of fixed energy which correspond to resonant waves about the nucleus, a lot like the resonant frequencies for the strings on a guitar

The Schrodinger wave equation HΨ = EΨ The main point is that Ψ can only be used to make statements as to the probability of locating the electron based on x, y and z-axes.

Probability density -

Probability density describes the probability of finding an electron at a point in space (from Ψ2 ) with respect to the distance, r, from the nucleus.

Radial density function -

Radial distribution function : probability of finding an electron (from Ψ2 ), but within a spherical shell - distance of r from the nucleus.

-

Since the volume of these shells increases as r increases, the function increases from zero through a maximum before tending to zero again.

Orbitals -

Each s-subshell only has one orbital. Which means each s-subshell can only hold a maximum of 2 electrons

-

They all have a spherical shape

-

Each p subshell holds three orbitals: px , py , pz . Which means each p subshell can only hold a maximum of 6 electrons

-

They all have a ‘dumbell’ shape

-

Each d subshell holds five orbitals: dxy, dyz, dxz , dx²-y² , dz² . Which means each d subshell can only hold a maximum of 10 electrons

-

They all have a ‘double-dumbell’ shape

Nodes A nodal plane represents a plane in space about which the wavefunction for a particular orbital has a zero probability amplitude.

s orbitals have n-1 spherical nodes where the wavefunction has a zero probability amplitude. ie. regions where the probability of finding an electron is zero

Week 4: Valence Bond Theory Weekly objectives • • • • • •

Discuss some of the important properties of covalent bonds. Discover correlations between electronegativity, polarity, bond lengths and bond energies. Introduce Valence Bond Theory. Define sp, sp2, sp3, sp3d & sp3d2 hybrid orbitals. Investigate how the different modes of orbital overlap manifest to give σ and π bonds. Discuss several examples of the hybridisation schemes of small organic molecules

Bond energies

Bond energies are defined as the amount of energy that must be supplied to break a particular chemical bond. Bond energies, like bond lengths, vary in ways that can be traced to atomic properties, and there are three consistent trends in bond energies: 1. Bond energies increase as more electrons are shared between the atoms. Shared electrons are a kind of “molecular glue”, so sharing more electrons strengthens the bond. 2. Bond energies increase as the electronegativity difference (Δχ) between bonded atoms increases. Polar bonds gain stability from the electrostatic attraction between the negative and positive partial charges around the bonded atoms. Bonds between oxygen and other second-row elements exemplify this trend.

3. Bond energies decrease as bonds become longer. As atoms become larger, the electron density of a bond is spread over a wider region. This decreases the net attraction between the electrons and the nuclei. The following table of bond energies illustrates this effect.

Dipole moments - if two or more atoms are bonded together, and they are not homonuclear, then the -

electrons forming these bonds will be unevenly distributed. This uneven distribution is caused by the unequal electronegativities of each atom. The asymmetrical distribution of electrons, called dipole moments, means that one end is slightly negative and the other is slightly positive. Dipole moment values are symbolised by the Greek letter mu, μ. As a general rule, the greater the difference in electronegativity, Δχ, the larger μ

Dipole moments are determined experimentally using a setup similar to the diagram on the right. When an electrical potential is applied across the plates, the molecule aligns spontaneously. The δ+ ends of the molecules point towards the negative plate, and the δ- ends point towards the positive plate. The measured dipole moment, μ, is determined on the alignment of these molecules.

Polarity A polar molecule is where all the dipole moments from each bond produce an unequal sharing of electrons across the entire molecule. One simple example is HF. It has a dipole moment between the H and F bond and this unequal sharing of electrons across the entire molecule makes the HF molecule polar. To draw the overall dipole moment of a molecule, an arrow (vector) starts from the δ- end of the molecule and points towards the δ+ end of the molecule

Sigma bonds A sigma bond is defined by location of the bonding electrons being between the nuclei as shown in the diagram below

Hybrid bonds sp bonds -

Mathematically combining the px and s- atomic orbitals Energy of the two hybrid orbitals is an average of the energy of the two atomic orbitals used Energy of the py & pz atomic orbitals remains unchanged

sp2 bonds -

Combining the px , py & s-atomic orbitals generates three, 3- fold sp2 hybrid orbitals pz atomic orbital remains unchanged

sp3 bonds -

Combining the px , py , pz & s-atomic orbitals generates these four hybrid orbitals No unhybridised s- or p- atomic orbitals remain

sp3d and sp3d2 bonds -

These hybridisation schemes additionally incorporate d atomic orbitals sp3d hybridisation corresponds to a trigonal bipyramidal geometry sp3d2 hybridisation corresponds to a octahedral geometry

pi bonds

Week 5: Molecular Orbital Theory Weekly objectives • • • • •

Discuss the principles of Molecular Orbital (MO) theory Identify MO’s that are σ or π, bonding or antibonding Construct Molecular Orbital diagrams for homonuclear diatomics H2 - Ne2 and write their electronic configurations Calculate bond orders, and relate them qualitatively to bond lengths and bond enthalpies Recognise why MO Theory is a superior model to VB Theory.

Molecular orbital theory -

Molecular Orbital Theory exploits the wave-like properties of electrons The model uses quantum-mechanical treatment of molecules similar to that used for isolated atoms, based on the Schrödinger wave equation Describes regions of space that an electron might occupy over a whole molecule Basic idea is that wavefunctions are constructed by taking linear combinations of the atomic orbitals of the constituent atoms Because the electrons are considered as wavefunctions, we need to consider constructive and destructive interference of the waves.

Molecular orbital theory is particularly useful for predicting whether molecules will exist given the constituent atom’s electron configuration, the strength of covalent bonds and the magnetic properties of molecules.

Co...