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60 Cards in this Set
- Front
- Back
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What does the oxidation number of an element tell you |
How many electrons an atom has accepted or donated to form an ion, or to form part of a compound |
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What are the rules for electron configuration |
. All Uncombined elements have an oxidation number of zero. The oxidation numbers of elements in a compound add up to zero. In a monoatomic ion ( an ion consisting of just one atom) the oxidation number is the same as the charge. For molecular ions the sum of the oxidation numbers equals the ions charge. There’s some oxidation numbers you just need to learn |
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What’s the oxidation number of oxygen |
Oxygen nearly always has an oxidation number of -2 except when it’s in peroxide’s (O2²- ) where it’s -1, and molecular oxygen (O2) where it’s 0. |
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What’s the oxidation number of hydrogen |
Always +1. Except when In metal hybrides (MHx) where it’s -1, and in molecular hydrogen (H2) where it’s zero. |
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What’s the oxidation number of chlorine |
Usually -1 |
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What’s the oxidation number of fluorine |
Always -1 |
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What’s the oxidation number of group 1 metals in compounds |
+1 |
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What’s the oxidation number of group 2 metals in compounds |
+2 |
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What’s the oxidation number of Al |
+3 |
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How do we represent the oxidation number for transition metals when it varies |
We use Roman numerals |
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What’s the oxidation number of iron in iron (II) Sulfate |
+2 |
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How does electronegativity correspond to assigning oxidation numbers |
Always assign the most electronegative elements oxidation number first if there is more than one non metal in a compound or ion. Electronegativity increases torwards the top right of the periodic table ( ignoring group 0) |
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What does it mean when an ion ends with ‘ate’ |
They contain oxygen and another elements |
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What does the oxidation number in an ion ending with ‘ate’ refer to |
The non oxygen atom |
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What’s an oxidising agent |
It accepts electrons and gets reduced |
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What’s a reducing agent |
It donates electrons and gets oxidised |
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Whys it called redox |
Because reduction and oxidation have to occur simultaneously |
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Identify the oxidising and reducing agents in this reaction: 4Fe + 302 —> 2Fe2 03 |
Iron has gone from having an oxidation number of 0 to +3. It’s lost electrons and has been oxidised. This makes it the reducing agent in this reaction. Oxygen has gone from having an oxidation number of 0 to an oxidation number of -2. It’s gained electrons and has been reduced. This means it’s the oxidising agent in this reaction. |
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Define oxidation |
Increase in oxidation number. Loss of electrons. Addition of oxygen |
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Define reduction |
Loss of oxygen. Fall in oxidation number. Gain of electrons |
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Define reduction |
Loss of oxygen. Fall in oxidation number. Gain of electrons |
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When is OILRIG used |
For ionic compounds |
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Show a half equation for oxidation of zinc |
Zn > Zn²+ + 2e^- |
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Show a half equation for oxidation of zinc |
Zn > Zn²+ + 2e^- |
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Show half equation for reduction of chlorine |
Cl2 + 2e- > 2cl- |
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Why is the oxidation number for cl on Cacl2 -1 and not -2 |
Because there are 2 Cls. Always work out the oxidation for one atom of the element |
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Explain how metals and non metals form compounds |
When metals form compounds they generally donate electrons to form positive ions meaning they usually have positive oxidation numbers. When non metals form compounds they generally gain electrons meaning gosh usually have negative oxidation numbers |
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Explain redox reactions of metals and acids |
Acid + metal > salt + hydrogen. Metal atoms are oxidised, losing electrons to form positive metal ions ( in salts). The hydrogen ions in solution are reduced, gaining electrons and forming hydrogen molecules |
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Hydrochloric acid plus magnesium? |
2HCl + Mg > MgCl2 + H2 Hydrogen is reducing because it’s oxidation is reducing from +1 to 0. Magnesium is oxidising because it’s oxidation number is increasing from 0 to +2. |
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What do you observe when an acid reacts with a metal |
effervescence because the H2 gas is evolved. The metal will dissolve and get smaller |
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Explain relationship between distance from nucleus and energy levels |
Shells further from the nucleus have a higher energy leveland a larger principle quantum number than shells closed to the nucleus |
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What’s the total number of electrons in principle energy levels 1,2,3 and 4 |
1 = 2. 2= 8. 3= 18. 4= 32 |
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What are energy levels split into |
Sub energy levels labelled s, p, d and f |
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How many electrons can each sub shell hold |
s = 2. P = 6. D = 10. F = 14 |
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What are orbitals |
Regions of space that can be occupied by a maximum of 2 electrons with opposite spin to each other |
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Explain how shells are split up |
Shells are spit up into shun shells which each has a different energy and have different numbers of orbitals which can each hold up to 2 electrons |
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What happens if there are two electrons in an orbital |
They have to ‘ spin ‘ in opposite directions. This is called spin pairing |
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What’s the shape of s orbital |
Spherical |
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What’s the shape of s orbital |
Spherical |
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What’s the shape of p orbitals |
Dumbbell shales . There are 3 p orbitals and they’re all at right angles do eachother |
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Show px orbital |
Back (Definition) |
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Show py orbital |
Back (Definition) |
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Show pz orbital |
Back (Definition) |
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What do orbitals represent |
The mathematical probability of finding an electron at any point within certain spatial distributions around the nucleus |
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What’s the order of filling |
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p |
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Show order of filling by electrons in boxes template |
Back (Definition) |
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What does the box represent on a spin diagram |
One orbital |
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What does the arrow represent on a spin diagram |
One electron |
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Why do the arrows go in opposite directions on a spin diagram |
To represent the different spins of the electrons in the orbital |
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How are orbitals of the same energy occupied |
Singly before pairing. Because this is more stable |
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What are the rules for electrons in boxes |
Electrons fill up the lowest energy sub shells first. Electrons fill orbitals with the same energy singly before they start sharing. |
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Explain 4s and 3d subjshells |
4s sub shell has a lower energy level than the 3d sub shell even though it’s principle quantum number is bigger. This means the 4s sub shell fills first |
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What are the exceptions to the 4s filling before 3d rule |
Chromium and copper add both more stable with one electron in their 4s sub shells. So their 4s sub shells remain partially filled as their 3d sub shells full up. Chromium = 1s²2s² 2p^6 3s²3p^6 3d^10 4s^1 |
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What does the highest energy orbital that’s filled determine |
The block of the periodic table that an element is found in. If the highest energy orbital that is filled is an s orbital then the element will be in the s block |
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Explain sub shell notation |
Big number is main energy level. Letter is name of sub shell. Small number is number of electrons in sub shell |
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Give an example of sub shell notation for oxygen |
8 electrons total. 1s² 2s² 2p^4 |
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Show the blocks of the periodic table |
Back (Definition) |
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Explain configuration of ions |
Do it the same way as an atom but is it’s positive charged remove the required electrons and if it’s negatively charged add the required electrons |
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Explain configuration of ions |
Do it the same way as an atom but is it’s positive charged remove the required electrons and if it’s negatively charged add the required electrons |
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Explain noble gas symbols |
Sometimes used in electron configuration. For example calcium ( 1s² 2s² 2p^6 3s² 3p^6 4s²) can be written as (Ar) 4s² because the non-outer shells have the same electron configuration as argon |
