# Elements of Life

A summary of the Elements of Life unit

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• Created by: R_Hall
• Created on: 13-12-12 19:47

## 1.1 Amount of Substance

• The link between the mass of an element and the number of atoms it contains is the relative atomic mass of an element
• The mole is the unit that measures the amount of a substance in a way that equal amounts of elements consist of equal numbers of atoms
• The mass of one mole (molar mass) is equal to relative atomic mass in grams
• Relative formula masses are used to compare substances
• In substances such as methane where the formula represents a discrete molecule, the RFM is often called the relative molecular mass
• The things that make up a substance are called formula units, they match the formulae of the substances
• The number of formula units in one mole of a substance is a constant. It is called the Avogadro constant and has a value of 6.02 x10
• The molecular formula tells you the actual numbers of the different types of atom, and the empirical formula tells you the simplest ratio of the numbers of different types of atom in a substance
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## 1.2 Balanced Equations

• Step 1- Decide what the reactants and products are
• Step 2- Write formulae for substances involved, incl state symbols
• Step 3- Balance the equation so that there are the same number of each atom on each side
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## 2.1 A Simple Model of the Atom

• The atomic number (Z) is the number of protons in the nucleus (equal to charge on the nucleus)
• The mass number (A) is the number of protons plus neutrons in the nucleus, If number of neutrons is N- A= Z+N
• Nuclear symbols identify mass number and atomic number as well as the symbol for the element
• Atoms of the same element which have different mass numbers are called isotopes. Have same proton number, electron number and charge, but different neutron number and mass number
• The relative atomic mass is an average of the masses of the isotopes (relative isotopic masses), taking into account their abundances
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## 2.2 Nuclear Reactions

• Some elements of some isotopes are unstable- their nuclei break spontaneously, and they are radioactive. They are called radioisotopes
• As they breakdown the emit rays and particles through radioactive decay. The 3 types of emission (alpha, beta and gamma) are able to knock electrons out of atoms (ionise)- ionising radiation
• In a nuclear fusion reaction, two light atomic nuclei fuse to make a single heavier nucleus of a new element, releasing lots of energy. Fusion reactions only occur at high temps where the nuclei are close together (lots of energy and attractive forces)
• At high temps in gas clouds of stars, the e- have energy to escape from nuclei. The gases exist in ionised form- plasma
• The time taken for half of the nuclei of a sample of a radioactive element is the half-life. For any given isotope, the half-life is fixed
• Some types of minerals in rock types contain radioactive elements, so the ages can be estimated.
• Carbon-14 and Carbon-12 exist in all living things and while living the ratio doesn't change. However when dead, there becomes more C-12 and less C-14. Can be used in archaeology to estimate age
• Radioisotopes can be injected into the body, to locate problems (eg. cancerous tumours). Must have a short h-l to allow detection but minimise harm
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## 2.3 Electronic Structure: Shells

• Electron shells are labelled by giving each a principal quantum number, n
• 1st shell, n=1. 2nd shell, n=2. Further away from the nucleus, bigger number- more energy
• A shell which contains max number of electron is a filled shell
• A useful model of the atom shows that it is composed of a core of a nucleus and inner electron shells, surrounded by an outer shell
• The number of outer shell electrons determines the group number
• The number of the shell which is being filled determines the period of the element
• First ionisation enthalpy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms
• The smaller the number of shells, the lower the ionisation enthalpy value (easy to ionise and very reactive
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## 3.1 Chemical Bonding

• Positively charged ions- cations, negatively charged ions- anions
• Ionic bonding- where a metal reacts with a non-metal. The cations and anions which are formed are held together by a strong electrostatic bond; the oppositely charged ions attract each other strongly- ionic bond. In solids, the ions form a giant lattice due to bonds. Ionic compounds have high melting points, dissolve in water and conduct electricity (molten or dissolved in water)
• Covalent bonding-  bonds formed by sharing of electrons. Atoms held together as +ively charged nuclei attracted to -ively charged electron. Electron pairs which form bonds- bonding pairs, pairs not involved- lone pairs. 1 bonding pair- single bond
• Covalent network structures have very high melting points, insoluble and conduct. Simple covalent's have low melting pt, do not conduct and are insoluble
• Dative covalent bonds- in a dative bond, both bonding electrons come from the same atom
• Polar bonds- Different atoms attract bonding electrons unequally. One atom gets a slightly -ive charge (greater share of bonding electrons). The electron pulling power is electronegativity
• Metallic bonding- Metal's lose outer e- to form a lattice of +ive ions with a sea of -ive delocalised electrons. Metals have high melting pt, insoluble and conduct (s or l)
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## 3.2 The Shapes of Molecules

• Atoms in a molecule arrange themselves so pairs of electrons (bonding and lone) are far away from each other as possible
• 2 groups of electron-s- linear shape, bond angle 180
• 3 groups of electrons- planar triangular shape, bond angle 120
• 4 groups of electrons- tetrahedral shape, bond angle 109
• 5 groups of electrons- triagonal bipyramidal shape, bond angle 90/120
• 6 groups of electrons- octahedral shape, bond angle
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## 6.1 Light and Electrons

• 2 light models- wave theory and particle theory
• Speed of light = wavelength x frequency
• Particle theory- light as a stream of packet of energy called photons. Energy of photon= frequency x Planck constant
• Atoms become excited by absorbing energy. When excited atoms lose energy and return to their ground state energy is emitted as electromagnetic radiation. The light can be split up into an emission spectrum
• An emission spectrum consists of a series of coloured lines on a black background
• An absorption spectrum shows black lines on a continuous spectrum background
• Lines get closer together when frequency increases, and the lines in both spectra of an element are in the same places
• When an atom is excited, electrons jump into higher energy levels (absorption). Later they drop back down (emission)
• An electron can only possess definite quantities of energy (quanta)
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## 6.5 Mass Spectometry

• Mass spectrometry is used to measure the atomic or molecular mass
• Time of flight mass spectrometers identify ions by measuring the time that sample ions take to fly a known distance in a constant electrical field
• Ions are produced as separate pulses from the sample inlet. They are then accelerated and pass through the flight path and into the ion detector.
• Heavier ions move more slowly through the flight path than lighter ions
• Kinetic energy= 1/2 mv2 (where m is mass and v is velocity)
• The detector produces a varying electrical current when hit by the ions, and computer converts into a mass spectrum, showing mass and relative abundance
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## 11.1 Periodicity

• The modern periodic table is  based upon one proposed by Mendeleev in 1869. He placed elements with similar properties underneath each other in columns, leaving gaps to be filled later. The elements were arranged in order of increasing relative atomic mass
• The modern periodic table is organised by increasing proton number. It is split up into 4 blocks (s, p, d and f) and the rows are periods and the columns are groups
• The occurrence of periodic patterns is called periodicity
• Density- mass per unit volume. Density increases across the period until group 3, then it falls
• When elements are boiled/melted the bonds between particles must be overcome. Both increase across the period until you get to group 4, where it starts to fall
• Atoms become smaller when you cross a period (despite there being more electrons and protons). This is because the  more oppositely charged protons and electrons there are, the stronger the forces between them, causing the electrons to be pulled closer to the nucleus
• The amount of energy needed to remove an electron from an atom is its ionisation enthalpy
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## 11.2 The s Block: Groups 1 and 2

• Group 1- alkali metals, Group 2- alkaline earth metals
• Elements become more metallic and reactive as you move down the group, and less metallic as you move across the period (group 1 more reactive than group 2)
• S block metals tend to be soft and weak, with low melting points, but pure metals are too reactive with water and oxygen to be of use or found in natural state (naturally commonly found in compounds)
• All elements react with water to form hydrogen and a hydroxide, become more reactive as you go down the group. M(group 2 metal)+ 2H2O -> M(OH)2 + H2
• General formula of oxides in MO and hydroxides is M(OH)2. In water, oxides and hydroxides form alkaline solution, although not very soluble- most strong at bottom of group
• Heating of carbonates= MCO3 -> MO + CO2. Carbonates become more difficult to decomposes you go down the group, lower elements have a higher thermal stability
• Hydroxides become more soluble as you go down the group
• Carbonates become less soluble as you go down the group- not a perfect trend though
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