Unit 6: Chemical Bonding (Test Review)!!!!:):):):):):)
Electron Dot Structures
A way to show the number of valence electrons in an atom! The number of valence electrons can be determined by the group the elements are in. These come from the s and p blocks only!
Metal Alloys
Alloy: mixture of two or more element, at least one of which is a metal! Properties are often superior to those of their component elements!
Polar bond
Bonds where the electrons are shared unequally between atoms! - The atom that is more electronegative will pull the electrons closer to itself! - When atoms are different, each has a different pull on the electrons if the E.N. difference > 0.4 Example: HF H - F The shared electrons are held closer to fluorine, because it is more electronegative. E.N. difference = 4.0 - 2.2 = 1.8!
Shape: Pyramid
Can only have 4 atoms! Central atom has one unshared pair! For the record, for all of these shapes, ALWAYS draw like the diagram to get credit! Degree: 107.5 degrees
Shape: Trigonal Planar
Central atom has NO unshared pairs Degree: 120 degrees Can only have 4 atoms!
Shape: Tetrahedral
Central atom has NO unshared pairs of electrons! Can only have 5 atoms! Degree: 109.5 degrees
Shape: Bent
Degree: 104.5 degrees Central atom can have one or two unshared pairs of electrons! Example: H2O Can have 3 atoms only!
General rules:
Different atoms around the central atom will always make polar molecules! Same atoms around a central atom always make non-polar molecules! Examples: H - H Non-polar. 2.2 2.2 = 0 (E.N. difference)! H - F 2.2 4.0 = 1.8 Polar!
Group 1 (electron activity)
Lose 1 Charge: +1
Group 2 (electron activity)
Lose 2 Charge: +2
Group 13 (electron activity)
Lose 3 Charge: +3
Group 14 (electron activity)
Lose/Gain 4 Charge: +/- 4
Covalent bonding
occurs between NONMETALS only!
Ionic bonds are
very strong! - Are solids at room temperature! - Have high melting and boiling points! - Made up of a metal and a nonmetal! - Does not conduct electricity in solid state, but does in the liquid (molten) and aqueous states!
Intermolecular Forces
Forces that attract molecules to OTHER molecules. These include: - Hydrogen bonding - Dipole-dipole attraction - London dispersion forces
Group 17 (electron activity)
Gain 1 Charge: -1
Group 16 (electron activity)
Gain 2 Charge: -2
Properties of Metals
Metals are good conductors of heat and electricity! Metals are malleable! Metals are ductile! Metals have high tensile strength! Metals have luster!
Formula Unit
The lowest/smallest whole number ratio of ions!
Valence electrons determine chemical reactivity
elements in the same group behave the same!
A single covalent bond is made up of
two electrons!
Group 17
7 valence electrons!
Group 18
8 valence electrons!
Polarity
A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment!
The Octet Rule
Atoms will gain or lose enough electrons to become isoelectronic with a noble gas!
Dipole-Dipole Attraction
Attraction between oppositely charged regions of neighboring molecules!
Properties of metals (day 6 MORE)
Malleable and ductile because sea of drifting electrons insulates the metal cations from one another! When subjected to pressure, cations easily move past each other like ball-bearings immersed in oil!
Metallic Bonding
Strong forces of attraction are responsible for the high melting point of most metals!
(structural formula) -
1 pair!
Hydrogen Bonding
Bonding between hydrogen and more electronegative neighboring atoms. Specifically Fluorine, Oxygen, and Nitrogen!
Shape: Linear
Can have 2 atoms: has only two balls Can have 3 atoms: has three balls Central atom must have NO unshared pairs! The degree of a linear shape is 180 degrees!
Covalent bonds are much
WEAKER than ionic bonds!
A crystal consists of
a 3D, repeating pattern of alternating + and - ions!
An ionic compound is made up of
crystals!
Aqueous
dissolved in water!
An ionic bond
forms from the attraction between + and - ions!
Hydrogen and halogens
will only form single bonds!
Group 15 (electron activity)
Gain 3 Charge: -3
Properties of Metals (day 6)
Good conductors because the electrons flow freely throughout the metal! As electrons enter one end of the metal the same number leave the other end!
If the electronegativity DIFFERENCE (all of the elements electronegativities subtracted) of the element/atom is less than 0.4 than it is nonpolar, if it is MORE or EQUAL to 0.4 than it is POLAR!
If the electronegativity DIFFERENCE (all of the elements electronegativities subtracted) of the element/atom is less than 0.4 than it is nonpolar, if it is MORE or EQUAL to 0.4 than it is POLAR!
Covalent Bonding
Involves a sharing of electrons. Atoms will share in order to reach a stable electron configuration!
Metallic Bonding (day 6)
The chemical bonding that results from the attraction between metal cations and the surrounding sea of electrons! Outer electrons are allowed to move freely throughout the metal! Valence electrons do not belong to any one atom!
When the electrons taken is shared unequally, like taking all the electrons from Oxygen and giving two to each Hydrogen and leaving Oxygen with no electrons, the bonds are polar. But the whole molecule itself, because it has all Hydrogens and no other element surrounding the central atom Oxygen, that makes the whole overall entire molecule NON-POLAR!!!!:):):):)
When the electrons taken is shared unequally, like taking all the electrons from Oxygen and giving two to each Hydrogen and leaving Oxygen with no electrons, the bonds are polar. But the whole molecule itself, because it has all Hydrogens and no other element surrounding the central atom Oxygen, that makes the whole overall entire molecule NON-POLAR!!!!:):):):)
One electron is donated to the bond from
each atom!
Ionic compounds have
high melting points! - Conduct electricity when melted or dissolved! Example: NaCL
The chemical formula for a covalently bonded group of atoms is called a
molecular formula! - Indicates the exact makeup of one molecule! - Example: C6H12O6
Atoms that are covalently bonded are called
molecules or molecular compounds!
To write the compound formula that would form between two elements in an electron dot structure,
write the metal first, then the nonmetal! Example: Na - F NaF! Remember, metals are on the left side and nonmetals are on the right side!
Group 1
1 valence electron!
1. Determine the total number of valence electrons in the molecule or ion. Add together the valence electrons from each atom. (Recall from Chapter 2 that the number of valence electrons is indicated by the position of the element in the periodic table.) 2. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond. In H2OH2O, for example, there is a bonding pair of electrons between oxygen and each hydrogen. 3. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen). These electrons will usually be lone pairs. 4. If any electrons are left over, place them on the central atom. Some atoms are able to accommodate more than eight electrons. 5. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or triple) bonds to the central atom to achieve an octet. This will not change the number of electrons on the terminal atoms!
1. Determine the total number of valence electrons in the molecule or ion. Add together the valence electrons from each atom. (Recall from Chapter 2 that the number of valence electrons is indicated by the position of the element in the periodic table.) 2. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond. In H2OH2O, for example, there is a bonding pair of electrons between oxygen and each hydrogen. 3. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen). These electrons will usually be lone pairs. 4. If any electrons are left over, place them on the central atom. Some atoms are able to accommodate more than eight electrons. 5. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or triple) bonds to the central atom to achieve an octet. This will not change the number of electrons on the terminal atoms!
- -
2 pairs!
Group 2
2 valence electrons!
- - -
3 pairs!
Group 13
3 valence electrons!
Group 14
4 valence electrons!
Group 15
5 valence electrons!
Group 16
6 valence electrons!
Non-polar bond
Electrons are shared equally between two atoms! - Atoms that are the same have the same pull on the shared electrons (same E.N. value or E.N. difference < or equal to 0.4)! Example: H2 H - H Both atoms have the same E.N. value! Example: CCL4 All atoms surrounding carbon are the same, even though the bonds are polar, they cancel each other out!
For example, chlorine, with seven valence electrons, is one electron short of an octet. If two chlorine atoms share their unpaired electrons by making a covalent bond and forming Cl2, they can each complete their valence shell! So like two elements with 6 valence electrons would bond to give each other 2 electrons and share them to reach an octet level of 8 valence electrons! So in a structural formula, the 4 dots circled would become 2 lines standing for double bonds! And in an atom with two different atoms bonding!
For example, chlorine, with seven valence electrons, is one electron short of an octet. If two chlorine atoms share their unpaired electrons by making a covalent bond and forming Cl2, they can each complete their valence shell! So like two elements with 6 valence electrons would bond to give each other 2 electrons and share them to reach an octet level of 8 valence electrons! So in a structural formula, the 4 dots circled would become 2 lines standing for double bonds! And in an atom with two different atoms bonding!
Structural formula example: H : H H . . H
H - H
Ionic Bonding
Involves a transfer of electrons. One element loses electrons and the other gains electrons!
O = C = S
Mixed bonds but it is polar molecule!
Covalent bonds have
Much lower melting and boiling points! - Is a solid, liquid, or gas at room temperature! - Non-conductors in any state!
Group 18 (electron activity)
No electron activity because it is a noble gas!
Mixed Bonds
Nonpolar and polar bonds!
H
Number of valence electrons: 1 Number of electrons short: 1 Number of bonds made: 1
Group 14 (BONDS)
Number of valence electrons: 4 Number of electrons short: 4 Number of bonds made: 4
Group 15 (BONDS)
Number of valence electrons: 5 Number of electrons short: 3 Number of bonds made: 3
Group 16 (BONDS)
Number of valence electrons: 6 Number of electrons short: 2 Number of bonds made: 2
Group 17 (BONDS)
Number of valence electrons: 7 Number of electrons short: 1 Number of bonds made: 1
O = C = O
Polar bonds but it is non-polar molecule!
H - N - H - H
Polar!
H - O - H
Polar!
Helpful tip for structural formulas:
Put the atom that wants to make the most bonds in the middle!!
Empirical Formulas
The chemical formula for an ionic compound is arranged in the smallest whole-number ratio is known as an empirical formula Example: Mg3N2 vs. Mg6N4
Valence Shell Electron Pair Repulsion Theory (VSEPR)
The electron pairs (both shared and unshared) in the outermost energy level try to get as far apart from each other as possible! - This determines the shape of the molecule!
The octet rule is a chemical rule of thumb that reflects the observation that elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. Helium thing of helium being in the noble gas row and having 8 valence electrons octet thing! ANYTIME ANY ELEMENT WANTS 8 VALENCE ELECTRONS LIKE HYDROGEN NEEDING JUST 1 MORE TO GET 2 TO BE LIKE HELIUM WHO'S ATOMIC NUMBER IS 2 MEANS THAT THEY ARE TRYING TO BE LIKE A NOBLE GAS NO MATTER HOW SMALL THE ATOMIC NUMBER IS ALL THAT MATTERS IS THE NOBLE GAS AND THAT IS IT! REVIEW THESE LARGE NOTES MANY TIMES!!!!!!!!!!!:):):):)
The octet rule is a chemical rule of thumb that reflects the observation that elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. Helium thing of helium being in the noble gas row and having 8 valence electrons octet thing! ANYTIME ANY ELEMENT WANTS 8 VALENCE ELECTRONS LIKE HYDROGEN NEEDING JUST 1 MORE TO GET 2 TO BE LIKE HELIUM WHO'S ATOMIC NUMBER IS 2 MEANS THAT THEY ARE TRYING TO BE LIKE A NOBLE GAS NO MATTER HOW SMALL THE ATOMIC NUMBER IS ALL THAT MATTERS IS THE NOBLE GAS AND THAT IS IT! REVIEW THESE LARGE NOTES MANY TIMES!!!!!!!!!!!:):):):)
Relative Magnitudes of Forces
The types of bonding forces vary in their strength: Covalent bonds (strongest) Hydrogen bonding (2nd strongest) Dipole-dipole interactions (3rd strongest) London forces (4th strongest)!
London (Dispersion) Forces
The weakest of intermolecular forces, these forces are proportional to the mass of the molecule! These forces occur due to temporary dipoles! Temporary dipoles are formed on molecules and atoms due to the random movement of electrons! These temporary dipoles cause induced dipoles of neighboring molecules! When the opposite poles of molecules meet they attract! These are the only forces of attraction between completely nonpolar molecules!
To transport electrons from one element to the other,
all the electrons from that element that it is being taken from have to go, so you could add more of the element that is taking them just in case you have more than two, which is what each element will take from that element, only TWO ELECTRONS!
