Atomic Structure

BI1014 Biological Chemistry

Pure substance

Pure substance

• matter that has a fixed or definite composition

• can be elements or chemical compounds

Element

simplest type of material made up of atoms

Compound

pure substance consisting of atoms of two or more elements always chemically combined in the same proportion

Mixtures

two or more different substances are physically mixed, but not chemically combined together

Homogenous mixture

composition is uniform throughout the sample

Heterogenous mixture

components do not have a uniform composition throughout the mixture

Conservation of Mass

mass cannot be created or destroyed; it simply changes forms.

Law of Definite Proportions

given chemical compound always contains the same elements in the exact same proportion by mass

Law of Multiple Proportions

whenever the same two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers

Atomic Structure

• Nucleus and shell

• Nucleus is positively charged and consists of protons and neutrons

• Nucleus contains 99.9% of mass

• Shell negatively charged contains electrons

• Shell 100,000 times larger than nucleus

Proton

• Symbol

• Charge

• Mass (amu)

• Location in atom

• p or p+

• + 1

• 1.007

• Nucleus

Neutron

• Symbol

• Charge

• Mass (amu)

• Location in atom

• n or n0

• 0

• 1.008

• Nucleus

Electron

• Symbol

• Charge

• Mass (amu)

• Location in atom

• e-

• - 1

• 0.00055

• Outside nucleus

Atomic number (Z)

number of protons in an atom which is equal to number of electrons

Mass number (A)

number of protons + number of neutrons

Isotopes

atoms of the same element that have the same atomic number but different number of neutrons

Avogadro’s Number

• 6.02 ×10²³

• Number of atoms in one mole of any element

• Number of atoms in 12g of ¹²C

Molar mass

quantity in grams that equals the atomic mass of that element/compound

Molar mass equation

Electronegativity

Electronegativity is a measure of the degree to which an atom ‘draws’ electrons towards itself in a bond

• The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionisation energy.

• Atoms with low ionisation energies have low electronegativities, because their nuclei do not exert a strong attractive force on the electrons, and vice versa.

• In a period, moving from left to right, electronegativity increases (as ionisation energy increases).

• Going down a group in the periodic table, electronegativity decreases as both atomic number and atomic radius increase, and ionisation energy decreases.

Ionisation energy

• Ionisation energy: energy required to completely remove an electron from a gaseous atom or ion

  • The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove.

  • Ionisation energies increase moving from left to right across a period (decreasing atomic radius), and decrease moving down a group (increasing atomic radius).

  • Group 1 elements have low ionisation energies, with the loss of an electron forming a stable octet.

Electron affinity

• Electron affinity: ability of an atom to accept an electron

  • Atoms with stronger effective nuclear charge have greater electron affinity.

  • For example, Group 17 elements, the halogens, have high electron affinities; the addition of an electron to such elements results in a completely filled outer electron energy level.

ED

Electron energy levels

Regions around an atom’s nucleus where electrons are found, with each level corresponding to a specific energy state.

• Main shells

  • Discrete energy levels

  • Main shells 1, 2, 3, 4, 5, 6, 7,

  • Each shell contains maximal 2 n² electrons (n = main shell number)

Subshells

• Main shells divided into subshells: s, p, d, f

• Each subshell contains a maximal number of electrons

  • S: 2 electrons

  • P: 6 electrons

  • D: 10 electrons

  • F: 14 electrons

• The higher the main shell, the more subshells

  • N = 1 only 1 s = 2 electrons

  • N = 2: one s and one p subshell = 2 + 6 = 8 electrons

  • N = 3: one s, one p and one d subshell = 2 + 6 + 10 = 18

Atomic orbitals

regions in space where there is a high probability of finding electrons

• Size reflects 95% probability of presence of electrons

• Number of nodes dependent on main shell n and defines kind of subshell

• Number of subshells reflect possible spatial orientation

• Atomic orbitals have particular shapes and can hold a maximum of two electrons. Closest to the nucleus is the lowest energy level with the value n = 1

Energy level of electrons

• Characterised by 4 quantum numbers

• Electrons differ in at least one number

  • Principle quantum number n (main shell): size of orbital(s)

  • Angular quantum number I (subshell): shape of orbital

  • Magnetic quantum number m (no of orbitals in subshell): orientation of orbital

  • Spin quantum number s

    • 'spin' of electron within orbital

    • Values -1/2 and + 1/2

    • Maximum of 2 electrons per orbitals

Group number

represents valence electrons

Period number

represents number of electron shells

Subshells in periodic table

Orbital filling order

• Electrons fill orbitals of the lowest energy first. So the first level (n = 1) with its 1s orbital is filled before an electron can occupy the second level.

• Within an energy level electrons fill the orbitals with the lowest energy first. s orbitals represent a lower energy configuration than p orbitals and so electrons will fill the s orbital within an energy level before filling the p orbitals.

• Individual orbitals can hold a maximum of two electrons. Each s orbital (1s, 2s, 3s etc.) can only accommodate two electrons at the most. When two electrons occupy the same orbital they spin in opposite directions, otherwise their negative charges would force them apart.

Electron configuration in boxes

• Each box represents an individual atomic orbital, and arrows (representing electrons) are used to show the specific placement of electrons within those orbitals.

• Electrons fill lower energy orbitals first, no two electrons in the same orbital can have the same spin and every orbital in a sub shell is singly occupied before any are double occupied

Summary of Trends in Periodic Properties of Representative Elements

• Period Properties - Valence electrons, Atomic size, Ionisation energy, Metallic character

  • top to bottom within a group

  • left to right across a period

• Valence electrons

  • Remains the same

  • Increases

• Atomic size

  • Increases due to the increase in number of energy levels

  • Decreases due to the increase of protons in the nucleus that pull electrons closer

• Ionisation energy

  • Decreases because valence electrons are easier to remove when they are farther from the nucleus

  • Increases as the attraction of the protons for outer electrons requires more energy to remove an electron

• Metallic character

  • Increases because valence electrons are easier to remove when they are farther from the nucleus

  • Decreases as the attraction of the protons makes it more difficult to remove an electron

Stable isotopes

Isotopes of an element that do not undergo radioactive decay and remain unchanged over time. Most naturally occurring isotopes of elements up to atomic number 19 have stable nuclei.

Radioactive isotopes

Elements with unstable nuclei due to uneven ratio of neutrons and protons which the nuclear forces cannot offset the repulsions between the protons. It spontaneously emits small particles called radiation to become more stable.

Alpha radiation features

• Emitted by heavy nuclei (elements)

• Helium nucleus

• 2 protons and 2 neutrons

• Mass number of 4

• Atomic number of 2

• Charge of 2+

• α symbol

Beta radiation features

• Emitted by light and heavy nuclei

• High energy electron

• Breakdown of a neutron into a proton and an electron

• Mass number of 0

• Atomic number of -1

• Charge of -1

• β symbol

Gamma radiation features

• Excited nuclei and excess energy

• Released when an unstable nucleus undergoes a rearrangement of its particles to give a more stable, lower energy nucleus

• Mass number of 0

• Atomic number of 0

• Charge of 0

• γ symbol

Position features

• Antimatter of an electron

• Proton in an unstable nucleus is converted to a neutron and positron. The neutron remains in the nucleus, but the positron is emitted from the nucleus.

• Mass number of 0

• Atomic number of 1

• Charge of 1+

• β ⁺ symbol

Orbital electron capture features

• One of the orbital electrons is captured by a proton in the nucleus, this creates a neutron and a neutrino which is emitted

• Electron is removed so atomic number decreases by 1

• ν _e symbol

Measures of radiation

• Activity: number of nuclear disintegrations per second - becquerel (Bq) which is 1 disintegration/s

  • Old unit: curie (Ci) number of disintegrations that occur in 1s for 1g of radium

• Absorbed radiation: amount of radiation absorbed by a gram of material - Gray 1J/kg = 1 Gy

  • Old unit: radiation absorbed dose 1 rad = 0.01 Gy

• Rem (radiation equivalent in humans): measures the biological damage - New unit: Sievert 1 Gy x factor = 1 Sv

  • Old unit: radiation equivalent in humans 1 rem = 1 rad x factor

  • Factors:

    • 1 for Beta, Gamma and X-rays

    • 10 for high-energy protons and neutrons

    • 20 for alpha radiation

Half life

• Amount of time it takes for one half of a sample to decay

Natural radioisotopes

Natural radioisotopes

Isotopes occurring in nature that are unstable and undergo radioactive decay. They have a long half-life

Artificial radioisotopes

Isotopes that are artificially created, are typically unstable and undergo radioactive decay, emitting radiation.

Applications of radioactivity

• Carbon dating

• Radio-immuno assay

• Medical

  • Visualisation

  • Treatment

• Energy generation

  • Fission

  • Fusion

Medical applications using radioactivity

• Radioisotopes go to specific sites in the body

• By detecting the radiation they emit, an evaluation can be made about the location and extent of an injury, disease, tumour, or the level of function of a particular organ

• Higher levels of radiation are used to treat or destroy tumours

Nuclear fission

The bombardment of a large nucleus breaks it apart into smaller nuclei, releasing one or more types of radiation and a great amount of energy.

Nuclear fusion

Small nuclei combine to form larger nuclei while great amounts of energy are released