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40 Cards in this Set
- Front
- Back
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State the relative mass and relative charge of a proton, a neutron, and an electron |
Proton: 1, +1 Neutron: 1, 0 Electron: 1/1840, -1 |
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Describe the structure of an atom |
A nucleus containing protons and neutrons, surrounded by energy levels called quantum shells, which contain electrons |
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Define Isotope |
Atoms of the same element which have the same atomic number but different mass numbers, due to a different number of neutrons |
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Define atomic number |
The number of protons in an atom (it is the bottom number when shown on the periodic table) |
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Define mass number |
The combined mass of protons and neutrons in an atom |
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State and explain whether or not isotopes react in the same way |
Isotopes of the same element would react in the same way as chemical reactions are to do with electrons rather than neutrons. Isotopes always have the same number of electons |
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Define relative atomic mass |
The weighted mean mass of an atom compared to 1/12 the mass of an atom of 12C |
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Define relative isotopic mass |
The mass of an individual atom of a particular isotope compared to 1/12 the mass of an atom of 12C |
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What does a mass spectrometer do? |
Measures the masses of atoms and molecules. Calculates the relative abundance of each isotope, displaying the information on a mass spectrum |
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What is the 'molecular ion peak' on a mass spectrum? |
The peak with the highest m/z ratio (careful, sometimes there is an m+1 peak) |
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What is the name given to the energy levels which electrons reside in? |
Quantum shells |
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How many subshells are there and what are they called? |
4: s, p, d, f |
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How many orbitals are there in a s subshell? How many electrons does it contain? |
1 orbital 2 electrons |
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How many orbitals are there in a p subshell? How many electrons does it contain? |
3 orbitals 6 electrons |
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How many orbitals are there in a d subshell? How many electrons does it contain? |
5 orbtials 10 electrons |
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How many orbitals are there in a f subshell? How many electrons does it contain? |
7 orbitals 14 electrons |
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Explain why electrons many be represented as arrows in boxes |
Electrons in each orbital have opposite spin: they "orbit" in opposite directions in the orbital, therefore they can be displayed in boxes according to each orbital which shows the direction of the electron in the orbital |
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Define orbital |
A region around the nucleus where there is a high probability of finding and electron, and which can hold up to two electrons of opposite spin |
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Describe the shape of an s orbital |
Spherical |
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Describe the shape of a p orbital |
A dumbbell |
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What are the vertical columns of the periodic table known as? |
Groups |
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What are the horizontal rows of the periodic table known as? |
Periods |
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State Hund's rule |
Electrons occupy orbitals singly before pairing happens |
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State the Pauli exclusion principle |
Electrons in the same orbital must have opposite spins |
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What is the 1s notation |
It is how you write out the electron configuration of a atom e.g Calcium is 1s2,2s2,2p6,3s2,3p6,4s2 |
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What do they numbers before and after the letter mean in 1s notation? |
The number before the letter refers to the shell. The number after the letter refers to the number of electrons in the subshell. E.g 3s2 mean 2 electrons in the s subshell of the third shell |
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Why do electrons occupy the 4s subshell before the 3d subshell? |
Because the energy of the 4s subshell is less than that of the 3d subshell and electrons always want to be in the subshell with the lowest energy |
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What are the exceptions to the rule where the 4s subshell is occupied before the 3d subshell? |
Copper (Cu) and Chromium (Cr) Cu: 1s2,2s2,2p6,3s2,3p6,3d10,4s1 Cr: 1s2,2s2,2p6,3s2,3p6,3d5,4s1 This makes them more stable |
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In terms of groups, where is the S block in the periodic table? |
Groups 1 and 2 |
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Where is the p block on the periodic table? |
Groups 3,4,5,6,7,0 |
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Give the general equation for first ionisation energy |
X(g) ----> X+(g) + e- |
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Define first ionisation energy |
The energy per mole taken to remove the outermost electron from 1 mole of gaseous atoms of an element |
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Write the general equation for second ionisation energy |
X+(g) ----> X2+(g) + e- |
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Describe the trend in successive ionisation energies and what these show evidence for |
The ionisations energy increases as more electrons are removed as there is less shielding and the electrons have stronger interactions with the nucleus. They provide evidence for quantum shells as there is larger increases in energy going from one shell to the next than from removing electrons from the same shell. |
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What is the trend in first ionisation energy going down a group? |
First IE decreases down a group as there are more quantums shells, therefore more shielding and outer electrons are further from the nucleus, thus have weaker interactions with the nucleus |
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What is the trend is first ionisation across a period? |
Generally increases, as atomic charge increase has a larger effect than increased electron repulsion. Exceptions to this trend are the 3rd and 6th elements in the period (e.g B and O in period 2) 3rd: because you are going from a p orbital to an s orbital. S electrons are more closely held to the nucleus so require more energy to remove 6th: Because hunds rule, pairing happens at the 6th element therefore more repulsion |
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What is atomic radius? |
The distance from the centre of the nucleus to the boundary of the electron cloud |
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Describe the trend in atomic radius across a period |
Atomic radius decreases as though you are adding an electron they are still in the same shell. However atomic charge increases so electrons are pulled closer (charge density increases) thus reducing atomic radius |
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Describe trend in atomic radius down a group |
Atomic radius increases down a group as the is greater shielding due to the addition of quantum shells and nuclear charge stays the same |
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What is periodicity? |
A repeating pattern of chemical and physical properties across different periods |
