Chemistry 1611 Notes SEM 1 PDF

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Chemistry 1611NotesYear 1, Semester 11. Atomic Structure1. list the particles that make up atoms, their symbols and their relative masses and charges calculate the average atomic mass from isotope information be able to balance nuclear equations NucleogenesisNucleogenesis: origins of the elements. T...

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Chemistry 1611 Notes Year 1, Semester 1

1. Atomic Structure 1.1 -

list the particles that make up atoms, their symbols and their relative masses and charges calculate the average atomic mass from isotope information be able to balance nuclear equations

Nucleogenesis Nucleogenesis: origins of the elements. There are 4 sub atomic particles:

Forces in an Atom: Nucleons are held together by the strong nuclear force (SNF); Force of attraction between nucleons. -

Only effective over short distances Strong enough to overcome the electrostatic force of repulsion between protons in nucleus.

Isotopes and average mass The atomic mass of an element is the weighted average of the atomic masses of each of the naturally occurring isotopes. Example:

Origin of elements and Radioactivity: Elements come from a series of nuclear reactions starting with Hydrogen in stars. These reactions go on to occur with heavier nuclei which fuse together to form increasingly larger atoms. Eg:

Nucleogenesis (nuclear reactions) produce nuclides that can be stable or unstable → unstable nuclei decay, releasing with kinetic energy of gamma radiation. These high energy products are known as radioactivity. Application: Radioactivity is utilised in PET scans (Positron Emission Tomography), which involves the collision of a positron and electron. Upon collision, gamma rays are formed at 180 degrees, which are then detected by the PET and thus provides an overall image/scan of organs such as the brain. (INSERT WORKSHEET L1)

1.2 -

Recognise the historical context of the Bohr model of the atom. Be able convert between the wavelength, frequency and energy of light. Be able to calculate the energy of a hydrogen orbit. Be able to calculate the atomic emission spectrum of a hydrogen atom.

Structure of atoms- History: 1808 J Dalton → proposed the atomic theory, (which replaced the previous theory of infinitely divisible matter): -

All matter consists of atoms that can neither be created or destroyed. Atoms of one element cannot be converted into atoms of another element. Atoms of an element are identical and different from other elements’ atoms.

1897 JJ Thomson → proposed the sub-atomic particle named the ‘electron’, which was smaller than the atom. -

-

First discovered using a cathode ray tube [a vacuumed glass tube with a cathode and anode at opposite ends]. A cathode ray (beam of negatively charged particles/electron) occurred when a voltage was applied to the cathode. Disproved Daltons theory, as it showed atoms were divisible into sub-atomic particles.

1909 E Rutherford → proposed that atoms had a positive core named the ‘nucleus’, and that the rest was empty space. -

-

Gold-foil experiment: alpha decay (a helium atom) was directed at a thin gold sheet. Most of the alpha particles passed through, whilst some were deflected → this was unexpected. Due to repulsion positive nucleus and the positively charged alpha particles. Led to conclusion that atoms weren’t solid spheres, but mostly empty space with most of the mass being in the nucleus. Rutherford calculated the size of the nucleus using these deflection measurements.

1909 N Bohr → proposed that electrons occupied specific energy levels (orbits). -

When an electron moves from energy levels, it absorbed or emitted energy. From lower ‘n’ to higher ‘n’ = absorbed energy. From higher ‘n’ to lower ‘n’ = released energy. This energy corresponds to light of a specific energy/frequency.

Electromagnetic radiation:

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Wavelength,  , lambda: The distance between two adjacent identical points of the wave.

-

Frequency, , nu: The number of wave crests passing a given point per unit time.

-

EMR travels at the speed of light (c = 3.00 x 108 ms-1) Wavelength and frequency are related to the speed of light: c = 

-

All radiation has the same speed, c, but the energy can vary. The higher the frequency, the more rapidly the wave is oscillating and thus the higher the energy: E = h

h = 6.626 x 10-34 Js

E = energy for 1 photon

Example 1: A radio station transmits at a wavelength of 2.84m. Calculate the frequency. Calculate the energy. Calculate the energy for one mole of radiation. (ANS: 1.056 x 108 s-1 = 106MHz) (ANS: 7.00 x 10-26 J) (ANS: E = 7.00 x 10-26 J x 6.022 x 1023 mol-1 = 0.0421 J mol-1) NOTE: Another way of interpreting the energy of light is to find out how much energy for one mole of radiation: ➔ E = Joules of energy for one photon to be emitted x 6.022 x 1023 mol-1 (Avogadro’s constant) Atomic Emission Spectra: Each atom absorbs/releases light of certain energy when its electrons climb/fall from energy levels. This emitting light from the atom is quantised (i.e. released in discrete packets of energy called ‘photons’). -

Electrons absorb energy that exactly corresponds to gaps between orbits/energy levels and jumps to the shell (and same goes when falling from shells). Thus, the patterns of lines on the Atomic emission spectrum are unique to each element. AAS proves Bohrs model (only for the hydrogen atom) different wavelengths of light released (and thus colours observed).

Bohrs model of electrons occupying orbits of certain energies is supported using the following equation for energy associated with a particular orbit/energy level:

As

E =

Efinal - Einitial

then

E= - 2.18 x 10-18 J (1/n2f - 1/n2i) Z2

Example: Calculate the wavelength of light emitted when an electron moves from the n = 3 to the n = 2 orbit of a hydrogen atom.

 E = - 2.18 x 10-18 J (1 / 22 - 1 / 32) (1)2 = - 3.03 x 10-19 J (minus indicates light emitted) Now E = h and E = hc /  So  = hc / E = (6.626 x 10-34 Js) (3.00 x 108 ms-1) / (3.03 x 10-19 J) = 6.56 x 10-7 m or 656 nm (this is the red light observed on H emission spectrum) Application: Na+ is the major ion in extracellular fluid. The body requires 1-2 mmol/day of Na+ while a typical daily diet contains 100 times this amount. The excess is excreted by the kidneys. Commonly, atomic absorption spectroscopy is used to determine sodium ion concentration in which the line in the spectra used for this measurement is at 589 nm

(INSERT WORKSHEET L2)

1.3 -

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Explain the meaning of the orbital quantum numbers, n l ml ms. Understand the designation of orbitals such as 1s, 3d, 4p, 4f. Recognise the shapes of s, p and d atomic orbitals. Determine the number of electrons in an orbital/sub-shell/shell.

Wave-mechanical model: Most atomic spectra are much more complex than expected from a Bohr model of electron arrangements: Thus, Bohrs model only works for H atoms or atoms with 1 electron (ions), as others have lines that it cannot explain. Quantum mechanics can explain these lines (De Broglie). The problem with Bohr’s model was that it said electrons were particles that existed in defined orbits: ➔ De Broglie disproved this and proved the wave nature of electrons, and hence the dual nature of particles and light. His equation relates wavelength to momentum:

 = h/mv Dual nature of electrons (a type of particle) are see through experimental observations: Wave Maltese cross – cathode ray cast a shadow (a property of waves, such as light)

Particle Paddle wheel - cathode ray caused the paddle wheel to move (hence it must have a mass) Magnetic/electric field deflection – cathode rays were deflected by these fields (hence it must have a charge, meaning it’s a particle)

Due to De Broglie’s’ proposal, this led to Schrodinger’s equation, proposing that electrons behaviour could be explained by treated them as matter waves (as opposed to particles) → quantum mechanical model. ➔ They were standing waves (points with 0 displacement nodes) with certain integer wavelengths. ➔ His equation is a model that shows electron’s as matter waves/wave functions with certain energies: Ĥ

= E

( = wave function)

Heisenberg’s uncertainty principal suggests that we cannot know both where an electron is ( x) and its speed ( v) at the same time (fuzziness): x v ≥ h/4m This quantum mechanical model (Schrodinger’s proposal) can only be solved if various boundary conditions are applied. That is, the waves must be standing waves that are: -

continuous single valued multiples of a whole number or half wavelengths

These waves surround the nucleus instead of having rings of electrons around the nucleus. This is because we never truly know where the electrons are around the nucleus → only determine the probability of its location. There are then discrete solutions that represent the energy of each electron orbital. The orbitals are described by quantum numbers.

Quantum numbers: Orbital: the region of space in which the electron is most likely to be found (quantum mechanics): Energy

-

Different shapes for these orbitals (s, p, d, f) Maximum of 2 electrons in 1 orbital

E=0

n=3

1st quantum number- Principal quantum number ‘n’ n=2

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Describes the size of the orbital (shell). n = 1, 2, 3, … As n increase, E of electron increases, resulting in a larger orbital size. The lower the n value, the more stable the orbital is due to attraction to nucleus.

n =1

2nd quantum number – Angular momentum quantum number ‘l’ -

Describes the shape of the orbital (referred to as sub-shells).

-

l = 0, 1, 2, … (n-1) l=0 l=1 l=2 l=3

s orbital p orbital d orbital f orbital

d orbital

p orbital

s orbital

3rd quantum number – Magnetic quantum number ‘ml’ -

Describes the orientation of the orbital (on the x, y, z planes). l = 0; ml = 0 → 1 orientation (s orbital) l = 1; ml = -1, 0, +1 → 3 orientations (px, py, pz orbitals l = 2; ml = -2, -1, 0, +1, +2 → 5 orientations (dx, dy, dz orbitals

4th quantum number – Spin quantum number ‘ms’ -

Describes the spin of the electron. Each orbital, uniquely described by n, l and ml, may contain a maximum of two electrons, one spin +1/2, and the other spin -1/2. This means that electrons behave in two ways when placed in an external magnetic field. This obeys Pauli’s exclusion principal which states no 2 electrons can have the same 4 quantum numbers. Energy ml = -2,-1,0,+1,+2 l=2 10e l =1

ml = -1,0,+1

6e-

l=0

ml = 0

2e-

l =1

ml = -1,0,+1

6e-

n=2

l=0

ml = 0

2e-

n =1

l=0

ml = 0

2e-

n=3

18e -

8e-

Applications:

Shell

Sub-shell

Orbital

Electrons

The modern periodic table! Explained later. The periodic table is displayed in such a way to represent the electron configuration of atoms using these s, p, d, f orbitals by being designed in ‘blocks’. (INSERT WORKSHEET L3 + TUTORIAL SHEET 1 + HOMEWORK SHEET)

2. The Periodic Table

2.1 -

be able to write out the electron configuration for atoms and ions. recognise ground state and excited state electron configurations. Be able to use the box notation to represent orbital occupancy. Understand the three rules associated with determining electron configurations.

Sub-shell energy: In Hydrogen, the energy of the sub-shells of a given ‘n’ are degenerate (of the same energy). In all other atoms the subshells are of different energies. The order in which subshells are filled is dependent on the energy of the orbital, as Aufbau’s principal states that the lowest energy orbital must be filled first (which depends on the value of ‘n’). 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p … [NOTE: look at periodic table → saves memorising]

Electron configuration: There are 3 rules used in determining electron configuration: 1. Pauli Exclusion Principle: No two electrons can have an identical set of four quantum numbers. I.e. there are a maximum of 2 electrons in any one orbital. A direct consequence of this rule is that any orbital can contain a maximum of 2 electrons, and that the two electrons in a full orbital must be of opposite spin. 2. AufBau's Principle: fill up lowest energy orbitals before high energy ones. Exception is d4 or d9 transition metals – Cr and Cu. Aufbau’s principal states that complete orbitals are most stable, whilst halffilled stables are second most stable. Hence Cr and Cu would exist as: Cr expect: [Ar] 4s2 3d4;

observe [Ar] 4s1 3d5

Cu expect:[Ar] 4s2 3d9;

observe [Ar] 4s13d10

Orbital capacities are as follows: → s has 2 electrons → p has 6 electrons → d has 10 electrons → f has 14 electrons 3. Hund's Rule: Orbitals with the same energy (i.e. the same sub-shell) have the maximum number of unpaired electrons. This rule summarises the way in which electrons occupy orbitals of equal energies. The lowest energy configuration involving orbitals of equal energies is the one with the maximum number of electrons with the same spin orientation. In other words, this means that electrons of the same spin orientation will be of a lower energy state than if they weren't. Ground state: the atom has no ‘excited electrons’. Examples of electron configurations:

Elements in the same group have the same valence electron configurations and, as a result, very similar chemical properties.

Box representation of orbitals: A box represents an orbital and an arrow represents an electron. Indicates occupancy of orbitals.

Ions: Forming a cation: remove electrons from the sub-shell with the highest principle quantum number (n) first. Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6 or [Ar] 4s2 3d 6

Eg:

Fe2+: 1s2 2s2 2p6 3s2 3p6 4s0 3d 6 or [Ar] 3d 6 Fe3+: 1s2 2s2 2p6 3s2 3p6 4s0 3d 5 or [Ar] 3d 5 ➔ So for period 4 elements this means the 4s electrons are removed first. Forming an anion; simply add an electron to the orbital currently being filled.

Excited configurations: Electrons can be excited to higher energy levels -

Ground state Na has 1s2 2s2 2p6 3s1 One excited state of Na has 1s2 2s2 2p6 3s0 3p1 Excited state relaxes to ground state emitting yellow light (as used in sodium street lights):

Na emission spectra

Application: Mendeleev’s periodic table represents the chemical properties of the elements → elements that have similar/same valence shell configuration were placed into the same groups (dependent on its s, p, d and f orbitals), hence allowing for chemical properties to be consistent within a group. (INSERT WORKSHEET L4)

sodium line

2.2 -

be able to give examples of periodic trends and chemical properties used to construct the Periodic Table. explain periodic trends in atomic radii and ionization energies in terms of size and effective nuclear charge. give examples of essential, toxic and medicinal elements and how the role of these elements relate to their chemical properties

Periodic trend: The chemistry of an element is directly related to its valence electron configuration. Many physical properties result from the attraction of the valence electrons to the nucleus of the atom. As atomic nucleus increases, no. protons and no. of electrons increases, attracting in a 1:1 ratio (steady increase as no. of protons = electrons). This attraction between an electron and proton depends on what shell; core electrons cause shielding to outer electrons in terms of nuclear attraction. ➔ Periodic Trends in ionisation energy, atomic radius and electronegativity result from the influence of effective nuclear charge. Ionisation energy: [NOTE]: Effective nuclear charge – The net positive charge experienced by a valence electron in a multielectron atom. This is impacted by; no. of electrons, no. of protons, and the shielding effect. Ionisation energy is the energy needed to remove an electron from its element. First ionization energy: M (g) → M+ (g) + eAcross a period: first ionization energy generally increases across a period because from left to right, the number of electrons increases, which causes greater attraction between the nucleus and the electron, making it difficult to remove. The increasing nuclear charge also drags the outer electrons closer to the nucleus. This reflects the effective nuclear charge experienced by the electron being removed. Down a group: first ionization energy decreases down a group, because as you go down the group, atomic radii increases, therefore, the attraction between the nucleus and the outermost electrons becomes less, which results in a lower ionization energy required to remove the outermost electrons.

= easier to remove this electron due to repulsion, hence why O has a lower I.E than N.

Atomic radius: Atomic Radius is defined as ‘half the distance between two atoms of the same element’. Across a period: atomic radius decreases across a period due to greater attraction to nucleus since there are more electrons attracting to protons (contracts) → since the number of protons is also increasing left to right, the effective nuclear charge increases across a period, causing this decrease. Down a group: atomic radius increases due to the addition of more shells down a group, where the valence electrons remain the same → since this is the same, these valence electrons are exposed to the same effective nuclear charge, but instead further from the nucleus, resulting in this increase. -

There is also the shielding effect in place, in which the valence electrons encounter less effective nuclear charge because they are simultaneously attracted to the positive nucleus and repelled by the negative charged electrons, thus increasing the atomic radii even more down a group.

Cation < neutral atom -

Cations have less valence electrons which are held by the same effective nuclear charge. Hence, less electrons have to share charge → stronger attraction to nucleus and thus being smaller

Anions > neutral atom -

Anions have more electrons which are held by the same effective nuclear charge. Hence, more electrons have to share charge → weaker attraction to nucleus and thus being bigger. [INSERT TO TRENDS IN IONIC RADIUS SHEET]

Electronegativity: Electronegativity is the ability of an atom to attract electrons of a covalent bond towards it. Across a period: Electronegativity increases due to decrease in atomic radii. Down a group: Electronegativity decreases due to increase in atomic radii, as the distance between the valence electron and the nucleus.

d-block elements: 4d and 5d elements (period 5 and 6) have very similar sizes (in terms of atomic radius) due to ‘lanthanide contraction’, which is the poor shielding effect of 4f (which occurs in the lanthanides, atomic numbers 5771).

Elements and the body + trace metals or toxic metals: Essential elements, or trace elements, are elements needed in the body in small amounts, which often include metals, as well as transitional metals (which have a wide range of oxidation states and form stable compounds). Toxic metals include any metals found in high enough concentrations (greater than the optimum range). Also, toxic metals can include those that are harmful in any concentration, and thus should not be in the body.

Category

Element

Essential

Na+

maintains osmotic balance and tissue water content of the body.

Small positive ion with low polarizability