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Unit 3: Chemical bonding (till page 31) Flashcards

1
Q

Valence shell is

A

the outermost shell

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2
Q

Valence electrons

A

outermost electrons
-important because they are involved in chemical reactions

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3
Q

Lewis Dot Structures

A
  • Show only the valence elections, ignoring the inner electrons because they are not involved in a chemical reaction.
    –> Making them more convenient & simpler than Bohr-Rutherford diagrams
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4
Q

In reality the octet rule is only true for…

A

the first 20 elements

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5
Q

Reactivity of Elements

A
  • Determined by # of valence electrons
    –> All atoms are stable when they have a full valence shell –> called stable octet. The closer the element is to having a stable octet, the more reactive it is. (alkali metals & halogens are most reactive)
    –> An atom will lose, gain, or share electrons to get a stable octet (similar to the nearest noble gas)
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6
Q

The main group metal atoms will lose ______ of their valence electrons

A

all

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7
Q

Type of atoms involved & type of chemical bond in ionic compounds

A
  • Metals & non-metals
  • ionic bond
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8
Q

Type of atoms involved & type of chemical bond in molecular compounds

A
  • non-metals
  • molecular bond (polar and non-polar)
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9
Q

Type of atoms involved & type of chemical bond in metallic compounds

A
  • Metals
  • metallic bond
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10
Q

What is an ionic bond?

A

The electrostatic (the forces that electric charges exert on each other) attraction between ions of opposite charges

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11
Q

What happens when we combine a metal & a non-metal together?

A
  • Metals have low IE (ionization energy–> the energy to remove an electron) and low EN (electronegativity –> ability to attract shared electrons in a chemical bond)
  • Non-metals have high IE & EN
  • The metal’s electrons will pulled closer to the non-metal’s nucleus than its own due to great EN gap (usually △EN > 1.7) –> therefore it will feel a greater attraction from non-metal nucleus than its own
  • As a result, metal loses its valence electrons –> positively charged (cation) & non-metal gains electrons –> negatively charged (anion)
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12
Q

How to calculate EN

A

Look at Electronegativity chart at back of periodic table
- find EN of elements in equation
- subtract smaller number from bigger
- Number of each atom doesn’t matter –> Ex. if you have 2 hydrogen atoms in the formula unit, do not multiply 2.1 by 2.

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13
Q

Naming Binary Ionic Compound

A

Metal + Non-Metal (ending with ide)

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14
Q

Writing Formula for Ionic Compounds

A

Criss Cross Method (be sure to simplify)

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15
Q

Binary ionic compound & Polyatomic Ionic Compounds

A

Binary: contains only 2 elements
Polyatomic: contains 3 or more elements

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16
Q

Crystal Lattice

A

These oppositely charged species being produced in close proximity are drawn together into an ordered, solid, three-dimensional array of cations & anions

17
Q

The smallest whole number cation-to-anion ratio in this structure represents the _____________

A

Formula Unit

18
Q

Solids tend to be very stable compounds. How does lattice energy explain the stability of ionic solids

A
  • the crystalline structure allows each ion to interact with the multiple oppositely charge ions, which causes a highly favourable (stabilizing) change in the enthalpy (sum of internal energy) of the system
  • A lot of energy is releases as the oppositely charged ions interact, causing the solids to be stable and have high melting and boiling points. Some compounds will decompose before they melt or boil
  • The vast number of interionic forces present in a crystal lattice, lock all the ions in place. (less movement, less energy, more stable). Helps explain why all ionic compounds are solids with high melting temperatures (ions are harder to pull apart)
19
Q

Lattice energy is

A

the energy released when gaseous ions of opposite charges are packed together to form an ionic solid
–> it is generally exothermic (energy is released) & has a major effect on whether a compound can be made

20
Q

Coulomb’s law to calculate lattice energy

A

F= k [(q1 x q2)/ r^2]
F: attractive force (lattice energy)
q: charge of ion
k: constant
r: distance between ions
(SEE PAGE 12CHAPTER 3: CHEMICAL BONDING BOOKLET

21
Q

Factors that effect the lattice energy of a compound

A

CHARGE FACTOR (q)
- Compounds with ions that have greater charges have higher lattice energy due to more electrons bein transferred, creating a stronger ionic bond
DISTANCE FACTOR (r)
- Compounds with ions that have smaller radii have larger lattice energies since the ions can get closer together, creating a stronger ionic bond
IONIC CHARGES HAVE A GREATER IMPACT ON LATTICE ENERGY (FORCE) THAN THE DISTANCE (RADII SIZE ) FACTOR
–> so when trying to see which ionic compound has a greater lattice energy, compare the charges first. If there is no difference, then check the radii sizes

22
Q

Lattice Formation - The Born Haber Cycle

A
  • The Born Haber cycle allows us to understand and determine the lattice energies of SOLID IONIC compounds
  • It is a cycle of energy change of process that leads to the formation of a solid crystalline ionic compound from the elemental atoms in the standard state. (regular state a room temp)
  • Since lattice energy is RELEASED when GASEOUS IONS bind to form an ionic solid, this process will always be exothermic, meaning the lattice energy value will be negative. It is normally expressed with the units kJ/mol.
23
Q

What happens when we combine a non-metal & a non-metal together?

A
  • The atoms will share electrons with each other to fulfill the octet rule arrangement –> each atom will have a full valence shell (same electron configuration as noble gases)
  • the bond is called covalent
  • The atom’s electron clouds will overlap as they are mostly empty space to cause attractive forces to exceed repulsive ones–> they cant be too close however, or the nucleus’ will repel
24
Q

Non-polar covalent bonds

A

If △EN (difference in electronegativity) is LESS THAN 0.5, the two atoms will have the same pull on the pair(s) of bonded electrons

25
A lone pair is
pairs of valence electrons that are not being shared
26
A bonding pair is
a pair electrons that are being shared by two atoms
27
Polar covalent bond
If △EN (difference in electronegativity) is GREATER or EQUAL to 0.5 and LESS than 1.7, the shared electrons will be drawn closer to the nucleus of the atom with the higher electronegativity - this unequal distribution of sharing of electron density will give that end of the bond a partially negative "pole" and the other a partially positive "pole". --> A bond "dipole" is said to exist and the bond itself is known as a polar covalent bond
28
Dipole moment
when you have a polar covalent bond, the electron density will be concentrated closer to to the atom with the greater EN, giving that end of the bond a partial negative charge and leaving the other atom end with a positive charge. Ex. In the ionic compound HCl, the bond dipole is said to be in the direction of Cl. (SEE PAGE 21 & 22 IN CHAPTER 3 BOOKLET TO SEE HOW TO LABEL PARTIAL CHARGES)
29
Naming Molecular Compounds
prefix+element 1 prefix+element 2 + ide ending - Prefix depends on number of atoms of each element in the compound. Do not use criss cross Prefixs: mono=1, di=2, tri=3, tetra=4, penta=5, hexa=6, hepta=7, octa=8, non=9, dec=10. - Cannot start with mono for the first element - Cannot have oo or ao together. Remove 1ST VOWEL
30
Properties of ionic compounds
- Are solids @ room temp and usually brittle (shatter easily) due to like-charge repulsion when ions are shifted (SEE PAGE 24 IN CHAPTER 3 BOOKLET) - the positive and negative ions are arranged in a crystal lattice - have high melting (>300 degrees Celsius) and boiling points due to strong electrostatic (the forces that electric charges exert on each other.) attraction (ionic bond) - As a solid, ionic compounds cannot conduct electricity, as the charged ions cant move --> however, when dissolved in water or in molten state (the charged ions are free to move), they can conduct electricity (electrolytes)
31
What conditions do compounds need to have in order to conduct electricity
1) charges (ions) --> full charges not partial 2) charges must be able to move around --> compound can't be solid
32
Properties of molecular compounds
- can be solid, liquid, or gas @ room temp. Small molecules are usually gases, large molecules are solids - usually have low melting & boiling points due weak attractions (intermolecular forces) between molecules. --> molecules are not bonded together so they are easier to pull apart rather than ionic formula units in crystal lattice - Do not conduct electricity when melted due to lack of charged ions--> they have a partial charge not a full charge - May dissolve (soluble) in water if they contain several N, O, of F atoms - Can contain several(100+) atoms - Can form strong covalent network solids (large molecular compounds held together by covalent bonds)
33
Metallic bonding
-Metallic bonds are present in pure elemental metals (only one element bonding with itself) , such as iron, or alloys (mixture of 2 or more metals) - metals lose their valence electrons (low EN & IE) M --> M^+ + free e^- - the electrons are relatively free to move from one positive nucleus to the next --> the positively charged metal ions are surrounded by a sea of moving electrons
34
What does the hardness & melting point of metallic bonding depend on?
- depends on how many valence electrons it has. --> The greater the number of valence electrons --> the more electrons the metals will lose--> the more positive the metal --> more free electrons and greater attraction between the electrons to the metal ions --> the more firmly the metal ions are held in place --> increase of hardness and melting point of the metal - use the F=k [(q1 x q2)/r^2] formula to compare hardness and melting points of metallic bonding
35
Properties of metallic bonds
- they are malleable (can be hammered into sheets) & ductile (draw out into thin wires) as metallic atoms can slide over each other easily - excellent conductors of heat (cuz moving electrons have KE) and electricity (meets both requirements) due to the free moving electrons
36
Basis of formation and relative strength of the intermolecular forces
Metallic bond: metal cations to delocalize electrons, strongest Ionic bond: cations to anions, 2nd strongest Polar covalent bond: partially charged cation to partially charged anion, 3rd strongest Nonpolar covalent bond: nuclei to shared electrons, 4th strongest

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