quizlet 2

• Atoms


  Made up of 3 subatomic particles - electrons, neutrons and protons
• Electrons


  - Subatomic particle

  - Charge of -1

  - Arranged in orbitals

  - Relative mass of 0.0005 - negligible
• Nucleus


  - Where most of the mass of the atom is contained

  - Made up of protons and neutrons

  - Diameter is much smaller than that of whole atom
• Neutrons


  - Subatomic particle

  - No charge

  - Relative mass of 1

  - Contained in nucleus

  - Dictates the isotope of an element that an atom is; not all atoms of the same element have the same number of neutrons
• Protons


  - Subatomic particle

  - Charge of +1

  - Relative mass of 1

  - Contained in nucleus

  - Dictates the type of element that an atom is; all atoms of the same element have the same number of protons
• Ions


  - Ions are charged atoms; positive ions have more protons than electrons, and vice versa for negative ions

  - Ions have different numbers of electrons to their parent elements'

  - e.g. Li⁺ has only 2 electrons, whereas Li has 3

  - e.g. F has 9 electrons, F⁻ has 10
• Isotopes


  - Isotopes of an element are atoms with the same number of protons but a different number of neutrons.

  - E.g. ³⁵Cl has 18 neutrons and ³⁷Cl has 20

  - Number and arrangement of electrons dictate the chemical properties of an element, so all isotopes of an element have the same chemical properties

  - However isotopes of the same element can have different physical properties such as density and diffusion rates
• Relative atomic mass


  The relative atomic mass is the weighted mean mass of an atom of an element compared to 1/12th of the mass of an atom of carbon-12

  
  

  Can be worked out from Isotopic Abundances - Multiply the isotopic mass of each isotope by its % abundance, add them up then divide the total by 100
• Relative isotopic mass


  The mass of an atom of an isotope of an element compared with 1/12th if the mass of an atom of carbon-12
• Relative molecular/formula mass


  The average mass of a molecule compared to the mass of an atom of carbon-12
• Mass Spectrometry


  Can be used to work out the relative atomic mass. Particles measured with a mass spectrometer must be charged, so they are often bombarded with electrons in order to remove one, giving a charge of +1

  
  

  1. Multiply each relative isotopic mass by its relative isotopic abundance, and add up the results

  2. Divide by the sum of the isotopic abundances

  
• Calculating Isotopic Masses from Relative atomic mass

  Need: Relative mass of element and all but one of the abundances and isotopic masses of its isotopes

  1. Find abundance of last isotope; percentage abundances so do 100-(sum of known% abundances)

  2. Put into equation for finding the relative atomic mass and rearrange for the unknown value
• Predicting mass spectra for diatomic molecules (E.g. Cl₂)


  1. Express each % as a decimal (e.g. 75%→0.75 and 25%→0.25)

  2. Make a table showing all the different Cl₂ molecules. For each, multiply the abundances of each isotope to get the relative abundance of each molecule.

  3. Look for any values in the table that are the same and add up their abundances

  4. Divide all the relative abundances by the smallest relative abundance to get the smallest whole number ratio. And by working out the relative molecular mass of each molecule, you can predict the mass spectra

  5. Plot the mass spectra with the relative abundances you worked out on the y-axis and the relative molecular masses (m/z) on the x-axis
• Identifying compounds using mass spectrometry

  1. Molecules in a sample are bombarded with electrons to remove an electron and form a molecular ion, M⁺

  2. The molecular mass is shown by the molecular ion peak - the peak with the highest m/z value, not including any M+1 peaks caused by presence of carbon-13
• Electron Shells


  - Made up of subshells and orbitals

  - Electrons move around the nucleus in quantum shells (aka energy levels)

  - Shells further from the nucleus have a greater energy level than those closer to the nucleus

  - Shells contain different types of subshell, each of which have different numbers of orbitals which can each hold 2 electrons
• Subshells


  This table shows the subshells and how many electrons can be contained in each.

  
• Orbitals


  - Orbitals within the same subshell have the same energy

  - s-orbitals are spherical

  - p-orbitals are dumbbell-shaped. There are 3 p-orbitals and they are at right angles to each other
• Electronic configuration


  - Electrons fill up the lowest energy subshells first

  - Electrons fill orbitals singly before they start pairing up

  - Exceptions: Chromium and Copper - donate a 4s electron to the 3d subshell because they are more stable with a full or half-full d-subshell
• Periodic table electron configuration blocks

  - s-block elements have an outer shell electronic configuration of s¹ or s²

  - p-block elements have an outer shell electronic configuration of s²p¹ to s²p⁶

  
• Atomic emission spectra - electron excitement

  - Electrons release energy in fixed amounts

  - In their ground state, atoms have their electrons in their lowest possible energy levels

  - If an atom's electrons take in energy from their surroundings, they can move up energy levels, getting further from the nucleus. These are known as excited electrons

  - Excited electrons release energy by dropping from a higher energy level down to a lower one. The energy levels all have fixed values - they are discrete

  - An emission spectrum shows the frequency of light emitted when electrons drop down from a higher energy level to a lower one. These frequencies appear as coloured lines on a dark background

  - Each element has a different electron arrangement, so the frequencies of radiation absorbed and released are emitted. This causes the spectrum for each element to be unique
• Atomic emission spectra


  - Each set of lines represents electrons moving to a different energy level

  - One set of lines is produced when electrons fall to the n=1 (ground state) level, another when they fall to n=2, etc.

  - When they drop to n=1, the series of lines is produced in the ultraviolet part of the electromagnetic spectrum

  - n=2 produces lines in the visible part of the spectrum

  - n=3 produces in the infrared part of the spectrum
• Emission Spectra support the idea of Quantum shells


  - Emission spectra show clear lines for different energy levels - supports idea that energy levels are discrete; electrons jump between levels with no in-between stage.
• Ionisation


  The removal of one or more electrons
• First Ionisation Energy


  The first ionisation energy is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous ions with a charge of +1.

  It is an endothermic process
• Factors affecting ionisation energy

  - Nuclear charge: more protons; more positively charged nucleus; stronger attraction for electrons

  - Electron shell: attraction falls off rapidly with distance; an electron in a shell close to the nucleus is much more strongly attracted than one in a shell further away

  - Shielding: as the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge and are repelled by the negatively charged electrons between them and the nucleus. This lessening of the pull of the nucleus by inner shells of electrons is called shielding.
• First ionisation energies decrease down a group


  As you go down the group in the periodic table, ionisation energies generally fall, i.e. it gets easier to remove outer electrons.

  This happens because elements further down the group have extra shells, so the atomic radius is larger, so the outer electrons are further away, which greatly reduces attraction.
• Successive ionisation energies


  - You can remove all the electrons from an atom, leaving just the nucleus

  - Each time an electron is removed, there's a successive ionisation energy which is greater than the previous ionisation energy

  - n'th ionisation energy can be written as:

  X(ⁿ⁻¹)⁺(g) →Xⁿ⁺(g) + e⁻
• Ionisation Energies show Shell structure

  - Within each shell, successive ionisation energies increase; less repulsion from other electrons each time, so there is stronger attraction to the nucleus

  - Big jumps occur between shells; an electron is being removed from a shell closer to the nucleus

  - This type of graph can tell you which group an element belongs to; count how many electrons are removed before the first big jump

  
• Modern period table organises elements by Proton Number

  - Dmitri Mendeleev created base for modern periodic table in 1869
• Electronic configuration decides the chemical properties of an element (periodicity)

  - s-block elements (Groups 1 and 2) have 1 or 2 outer shell electrons, which are easily lost to form positive ions with the electron configuration of an inert gas e.g. Na: 1s² 2s² 2p⁶ 3s¹ → Na⁺: 1s² 2s² 2p⁶ (electron configuration of Neon)

  - p-block elements (Groups 5,6,7) can gain 1,2 or 3 electrons to form negative ions with an inert gas configuration e.g. O: 1s² 2s² 2p⁴ → O²⁻: 1s² 2s² 2p⁶

  - groups 4-7 can also share electrons when they form covalent bonds

  - Group 0 (inert gases) have completely filled s and p subshells and don't need to gain, lose or share electrons; their full subshells are what make them inert

  - d-block elements (transition metals) tend to lose s and d electrons to form positive ions
• Atomic Radius decreases across a period (periodicity)


  This is because:

  - As the number of protons increases, the positive charge on the nucleus also increases, causing electrons to be pulled closer to the nucleus.

  - The extra electrons that the elements gain across a period are added to the outer energy level so they don't provide any extra shielding.
• Ionisation energy increases across a period (periodicity)

  This is due to:

  - number of protons increasing; stronger nuclear attraction

  - All extra electrons are at roughly the same energy level, even if the outer electrons are in different orbital types

  - Generally little extra shielding or extra distance to lessen nuclear attraction
• Drop in ionisation energy between groups 2 and 3 shows subshell structure


  This is because group 3 elements have their outermost electron in a p orbital rather than an s orbital; the outermost electron therefore is further away from the nucleus, this, and the shielding provided by the s orbital below the p orbital, is enough to override the effect of the increased nuclear attraction
• Drop in ionisation energy between groups 5 and 6 is due to electron repulsion


  - Elements with singly filled or full subshells are more stable than those with partially filled shells, hence they have higher first ionisation energies

  - e.g Sulphur 1s²2s²2p⁶3s²3p⁴ (1IE: 1000kJmol⁻¹ and Phosphorus 1s²2s²2p⁶3s²3p³ (1IE: 1012kJmol⁻¹)

  - shielding is identical in both, and electron being removed from same orbital

  - In P, electron is being removed from a singly-occupied orbital, in S its being removed from an orbital containing 2 electrons

  - The repulsion between two electrons in an orbital means that electrons are easier to remove from a shared orbital
• Bond strength affects melting and boiling points across a period


  - As you go across a period, the type of bond formed between atoms of an element changes

  - For metals, melting and boiling points increase across a period because the metallic bonds get stronger due to increased charge density

  - Elements that form giant covalent lattices (C and Si) have strong covalent bonds between all their atoms; a lot of energy is needed to break all these bonds; they have extremely high melting points

  - Simple molecular structures (e.g. N₂, O₂, P₄, etc) have low melting and boiling points due to the weak London forces between their molecules. More electrons means stronger London forces

  - Noble gases have the lowest melting points in their periods due to being monoatomic; they have very weak London forces
• Trends across Periods 2 and 3 for melting and boiling points


  Relatively strong increase across the period (metals→giant lattices) until group 5, where there is a sharp drop (due to simple molecular structures), then slight downward trend across the rest of the period.