1
Q
Atomic Structure
A
- Dalton - Solid Spheres
- Thomson - Electrons
- Rutherford - Nucleus
- Chadwick - Neutrons
- Bohr - Energy Levels
- Schrödinger - Electron Clouds
2
Q
Octet Rule
A
- Atoms are stable when the valence shell is full
- For many common elements they are stable with 8 electrons so this is generalized with the octet rule
- Atoms will form bonds or ions to achieve a complete valence shell
3
Q
Cations
A
+
- Cations form when an atom loses its electrons
- Metals form cations
- The energy required to “lose” an electron is called ionization energy
4
Q
Anions
A
-
- Anions form when electrons are gained by an atom
- Non-metals form anions
- The energy released with the electron is called electron affinity
5
Q
Isoelectronic
A
When atoms form ions, they often have a “noble gas-like” arrangement and atoms/ions end up with the same number of electrons. This is called isoelectronic.
6
Q
Isotopes
A
- A form of an element in which the atoms have the same number of protons but a different number of neutrons. So they have different masses
- You can determine the isotopic abundance of a sample using a mass spectrometer.
7
Q
Radioactive Decay
A
Radioactive Decay is the spontaneous disintegration of an unstable isotope
8
Q
Nuclear Radiation
A
Nuclear Radiation is energy or small particles emitted from a radioisotope as it decays
9
Q
Radioisotopes
A
Radioisotopes is an isotope that spontaneously decays to produce 2 or more smaller nuclei and radiation
10
Q
Radioactive
A
Radioactive means that a substance has the potential to emit nuclear radiation on decay
11
Q
Types of Radiation
A
Alpha Particle - a positive charged particle with the same structure as the helium particle
Beta Particle - a negative charged particle that is identical to an electron
Gamma Ray - a form of high-energy particle electromagnetic radiation emitted by some radioisotopes
12
Q
Equations
A
Check document
13
Q
Determining Atomic Mass of Elements
A
Atomic mass = (% abundance of isotope 1 x mass of isotope 1) + (% abundance of isotope 2 x mass of isotope 2) / 100
14
Q
Groups
A
- Elements show similar chemical properties
- Elements show similar trends in their chemical properties
15
Q
Periods
A
- As you move across periods, changes in the chemical and physical properties that are repeated in the next period
- This is what “period” and “periodic” refers to
16
Q
Physical Properties
A
Identify and explain the trends in the physical properties of the first 20 elements including: Atomic radius Ionic radius First ionization energy Electronegativity Melting point
17
Q
Atomic Radius
A
- This is the size of the atom
- There is no simple measure as atoms do not have a well defined “edge”
- We use the: covalent radius
- This is half the distance between the nuclei of two atoms in a covalent bond
- This means we do not have values for the noble gases as they do not form bonds
- The main factors influencing atomic radius are:
- Number of shells (the principal quantum number)
- The charge of the nucleus
18
Q
Periodic Trends - Atomic Radius
A
- Bottom-left - large atomic radius
- Top-right - small atomic radius
19
Q
Ionic Radius
A
- This is the “size” of an ion and is measured in a similar way to atomic radius
- The main factors influencing ionic radius are:
- Number of shells (the principal quantum number)… don’t forget this can be affected by the type of ion formed
- The charge of the nucleus
20
Q
Periodic Trends - Ionic Radius
A
- Positive ions
- Smaller than the ion they are derived from (Na bigger than Na+)
- Isoelectronic positive ions
- More protons, smaller radius, pull in electrons
- Negative ions
- Bigger than the ion they are derived from (Cl smaller than Cl-)
- Isoelectronic negative ions
- More protons, smaller radius, pulls in electrons
21
Q
First Ionization Energy
A
- This is the energy required to remove one mole of electrons from one mole of gaseous atoms to form positive ions
- Ex: Ag (g) → Ag+ (g) + e-
- Values are positive because this is an endothermic process
- Values are influenced by:
- Number of inner electron shells (and their shielding)
- Charge on the nucleus
22
Q
Periodic Trends - First Ionization Energy
A
- Lower ionization energy in lower left corner
- Higher ionization energy in upper right corner
23
Q
Electron Affinity
A
- Define as the enthalpy change when an electron is added to an isolated atom in the gaseous state
- Ex: X (g) +e- → X- (g) measured in KJ mol-1
24
Q
Periodic Trends - Electron Affinity
A
- Lower electron affinity in lower left corner
- Higher electron affinity in upper right corner
25
Electronegativity
- This is a measure of the degree to which an element attracts the shared oaur if electrons in a covalent bond
- Again, this means there are no values for the noble gases
- Values range over:
- 4.0 for Fluorine
- 0.7 for Francium
- Values are influenced by
- Number of inner electron shells (and their shielding)
- Charge on the nucleus
26
Periodic Trends - Electronegavity
- Lower electronegativity in lower left corner
| - Higher electronegativity in upper right corner
27
Melting Point
- This is the temperature (in Kelvin… i.e. Celsius +273) at which an element melts
- Values are influenced by:
- Nature of bonding: giant covalent, giant ionic metallic
- Strength of bonding
- Strength of intermolecular forces
28
How Ionic Compounds Form
- Metals lose their electrons to form positive ions
- Nonmetals gain electrons to form a negative ion
- The + ion and the - ion are held together by an electrostatic force. We call this an ionic bond
29
Ionic bond
Electrostatic attraction between two oppositely charged ions
30
Naming Binary Ionic
Compounds
| The name of the metal come first, non-metal ending changes to “ide”
31
Writing Formulas - Use the Criss-Cross
Look up charges on PT. Write them as superscripts, criss-cross (reduce if necessary) to get subscripts.
32
Naming Multivalent Metals
Need to use Roman Numerals to indicate which form of the metal is being used
33
Writing Formulas for Multivalent Metals
Use the charge given in the Roman Numerals and use the criss-cross rules
34
Hydrates
- Have water as part of its crystal structure
| - Provide the name of the ionic compound, then use prefixes from covalent compounds to indicate the number of “hydrate”
35
Polyatomic Ions
An ion that has more than one element bonded together
36
Important Polyatomics (Need to Know)
```
NO3 - → nitrate
SO4 2- → sulfate
OH - → hydroxide
PO4 3- → phosphate
CO3 2- → carbonate
NH 4+ → ammonium
```
37
Naming Ionic Compounds with Polyatomic Ions
- Name the metal (use Roman Numerals if necessary) and then the name of the polyatomic ion from the cart… don’t change the ending
- Note ammonium would come first because it’s positively charged and then the non-metals would change to -ide
38
Writing Formulas with Polyatomics
Put the polyatomic ion in brackets, use the criss-cross method. Leave the polyatomic in brackets if there is a subscript
39
Covalent Bonds Make Molecular Compounds
- When the nuclei of two atoms are both attracted to one or more pairs of shared electrons, the electrostatic attraction is called a covalent bond
40
Covalent bond
- Electrostatic attraction between the nuclei of two atoms and one/more shared pair(s) of electrons
41
Multiple Bonds
- Sometimes, atoms can share more than one pair of electrons between each other. This can result in double and triple bonds
42
Molecular Elements
- A pure substance composed of molecules made up of two or more atoms of the same element
- Most common are our diatomic elements (BrOHFINCl)
- Suffixes -gen and -ine are diatomic elements
43
Molecular Compounds
- Compounds made up of two or more nonmetals
| - H2O → 1 x 2 + 6 = 8 e-
44
Naming Molecular Compounds
- Use the prefixes (except we don’t use mono on the first element)
- No change to the first element and the ending of the second element changes to -ide
45
Writing Formulas
- The prefixes tell you the numbers (DO NOT CRISS-CROSS)
46
Rules for Drawing Lewis Structures
- Add up the valence electrons for all the elements involved
- If there is a - charge add the appropriate number of electrons
- If there is a + charge subtract the appropriate number of electrons
- Arrange the atoms evenly around the central atom
- Start with a single bond between each element
- Fill the octet of the outer elements (remember some just need 2)
- Then fill the central atom octet… you may need to use double or triple bonds!
- Just don’t exceed 8 electrons around each atom
47
So how do we determine if a substance contains covalent or ionic bonds?
- EN diff
- We need to calculate the electronegativity difference
- 0.0 → 0.4 = pure covalent bond
- 0.4 → 1.8 = polar covalent bond
- 1.8 → 4.0 = ionic bond
Level 3: Chemistry
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