Matter and Materials ATOMIC BONDS Grade 11 2018 Sutherland high school Mrs. Harrison
1. Chemical Bonds Why would atoms want to bond? Atoms are not generally found alone. They are found as components of larger units (molecules). Atoms strive to have the same electron structure as that of nobles gases. Noble gases have low energy levels. Why? Atoms, on their own, have a high potential energy. They want to bond with other atoms to form molecules that have a lower potential energy and are much more stable. More energy = less stable Less energy = more stable
2. Bonding Models COVALENT BONDS Between non-metals. Electrons are shared. Smallest particle is a molecule. Electronegativity difference less than 1.9 NON-POLAR COVALENT atoms attract electrons equally. POLAR COVALENT One atom attracts electrons more than the other.
2. Bonding Models IONIC BONDS Between metals and non-metals. Electrons are transferred. Smallest particle is an ion. Electronegativity difference more than 2.1 Atom 1 low ionisation energy. Atom 2 high electron affinity.
3. Valence Electrons The number of electrons in the outermost energy level of an atom. Corresponds to an elements group number on the Periodic Table. 4. Valency Number of electrons that will be shared, donated or transferred. Does not include charge.
5. Lewis Structure Illustrates the valence electrons of an atom. Electrons are drawn as dots around the symbol of that element. Complete activity 1
Paired electrons unpaired electrons
7. Single covalent bonds Different atoms that each have 1 unpaired valence electron will form a single bond. (share) Same atoms that each have 1 unpaired valence electron will form a single bond. (share)
7. Single covalent bonds Different atoms that each only have paired electrons will not bond. + no bond There needs to be incomplete or empty orbitals in order for a bond to take place.
8. Molecules with multiple bonds Atoms that share more than one pair of electrons to achieve noble gas structure.
The Octet Rule All atoms (except hydrogen and helium) want to have eight electrons, in four pairs of two, surrounding them in order to achieve noble gas structure.
Lewis Notation For more complex molecules Hand out worksheet PCl 3 OF 2 H 2 O CO 2 1. Central Atom. 2. Valence Electrons. 3. Bonding electrons 4. Lone pair electrons. NO 3 5. Ensure central atom has 8 surrounding electrons. NH 4 + H 3 O +
HOMEWORK EXERCISE 1 PG. 22-23
8. VSEPR Model Valence Shell Electron Pair Repulsion Model. Used to predict the shape of molecules. Molecules take on different shapes because the electron pairs repel each other. They are 3 dimensional and the atoms are arranged at specific angles between their pairs because of this repulsion. The shape of the molecules determines its chemical and physical properties.
THE 5 IDEAL MOLECULAR SHAPES ACCORDING TO THE VSEPR MODEL ARE:
worksheet
HOMEWORK EXERCISE 2 PG. 28
9. Electronegativity This is a measure of an atom s attraction on a shared pair of electrons. Atoms that have a high electronegativity strong attractive force gain electrons. Atoms that have a low electronegativity weak attractive force donate electrons. Electronegativity is influenced by: 1. Size of atom 2. Size of charge on nucleus
9. Electronegativity Electronegativity increases from left to right in a period because: The charge of the nucleus increases The size of the atom decreases There is a bigger attraction between the nucleus and the shared electron pair. The most electronegative element is fluorine. If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table.
9. Electronegativity difference This can be used to predict the type of bond that will form between the two atoms. Bond Non-polar covalent Covalent and very weak polar Polar covalent Ionic transfer of electrons occurs. Difference in electronegativity (ΔEN) 0 ΔEN < 0,4 0,4 < ΔEN < 1 1 < ΔEN < 2,1 ΔEN > 2,1 Noble gases have no value for electronegativity. They do not form (covalent) bonds. Atoms will attract the shared pair of electrons equally. One atom will attract the shared pair of electrons more than the other. One atom becomes slightly negative and the other slightly positive.
The distance of the electrons from the nucleus remains relatively constant in a periodic table row, but not in a periodic table column. The force between two charges is given by Coulomb s law. F=kQ1Q2r2F=kQ1Q2r2 In this expression, Q represents a charge, k represents a constant and r is the distance between the charges. When r = 2, then r 2 = 4. When r = 3, then r 2 = 9. When r = 4, then r 2 = 16. It is readily seen from these numbers that, as the distance between the charges increases, the force decreases very rapidly. This is called a quadratic change. The result of this change is that electronegativity increases from bottom to top in a column in the periodic table even though there are more protons in the elements at the bottom of the column. Elements at the top of a column have greater electronegativities than elements at the bottom of a given column. The overall trend for electronegativity in the periodic table is diagonal from the lower left corner to the upper right corner. Since the electronegativity of some of the important elements cannot be determined by these trends (they lie in the wrong diagonal), we have to memorize the following order of electronegativity for some of these common elements. F > O > Cl > N > Br > I > S > C > H > metals The most electronegative element is fluorine. If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table. Note: This simplification ignores the noble gases. Historically this is because they were believed not to form bonds - and if they do not form bonds, they cannot have an electronegativity value. Even now that we know that some of them do form bonds, data sources still do not quote electronegativity values for them. Trends in electronegativity across a period The positively charged protons in the nucleus attract the negatively charged electrons. As the number of protons in the nucleus increases, the electronegativity or attraction will increase. Therefore electronegativity increases from left to right in a row in the periodic table. This effect only holds true for a row in the periodic table because the attraction between charges falls off rapidly with distance. The chart shows electronegativities from sodium to chlorine (ignoring argon since it does not does not form bonds).
9. Electronegativity Determining whether covalent bonds are POLAR or NON POLAR Difference in electronegativity: EN of chlorine: 3.0 EN of hydrogen: -2.1.. EN = 0. 9 Cl H δ δ + Unequal spread of charge = Polar bond 0.4 2.1
9. Electronegativity Determining whether covalent bonds are POLAR or NON POLAR Difference in electronegativity: EN of chlorine: 3.0 EN of hydrogen: -3.0.. EN = 0. 0 δ Cl Cl δ Equal spread of charge = non-polar bond 0.4 2.1
9. Electronegativity Determining whether covalent bonds are POLAR or NON POLAR Difference in electronegativity: EN of oxygen: 3.5 EN of hydrogen: -2.1.. EN = 1. 4 H O δ H δ + H δ + Unequal spread of charge = polar molecule 0.4 2.1
9. Electronegativity Determining whether covalent bonds are POLAR or NON POLAR Difference in electronegativity: EN of oxygen: 3.5 EN of carbon: -2.5.. EN = 1. 0 δ δ + O C O δ Unequal load distributions on either side cancel each other out = non polar molecule 0.4 2.1
9. Electronegativity Determining whether covalent bonds are POLAR or NON POLAR Difference in electronegativity: EN of fluorine: 4.0 EN of boron: -2.0.. EN = 2. 0 δ δ δ + δ O Unequal load distributions on either sides cancel each other out. Symmetrical molecules = non polar molecule 0.4 2.1
Homework Exercise 3
10. Forces between atoms Force of attraction Force of repulsion Atom A Atom B
10. Forces between atoms A chemical bond will form when electrons are attracted simultaneously by two or more positive charges (nuclei). In a chemical bond, potential energy is stored. When bonds break or form, there are changes in potential energy.
11. Electrical PE changes in molecule formation Atoms always want to be stable. This means that they want the lowest possible energy. Bonds between atoms will form if the product has a lower potential energy than the atoms.
11. Electrical PE changes in molecule formation P : Atoms are far apart. PE is about 0. N: Atoms come closer together due to attractive forces. PE decreases. M: Atoms bond. Most stable position. Lowest PE. BM: Bond energy. Energy released when 2 atoms bond chemically. Always value (below x axis) because energy is absorbed. Higher BE, more difficult it is to separate the atoms in the molecule. OB: Bond length. Distance between the atoms (position) when the bond will form.
12. Bond energy and length Bond energy is the energy necessary to break a bond. The stronger the bond, the more energy you will need to break it. Bond energy is influenced by: 1.Length of the bond 2.Size of the atom 3.Bond order Longer bond weaker bond Bigger atoms longer bonds weaker less energy Order 1 (single) order 2 (double) order 3 (triple)
12. Bond energy and length Bond length is the distance between the two nuclei of the bonded atoms. Since the valence orbitals must overlap for a bond to form, the distance between the nuclei will be less than the sum of the radii of the two individual atoms. Bond length is influenced by: 1.Radii of individual atoms 2.Bond order Larger atom longer bond Higher order shorter bond length
Using Bond energies and Lengths What do you notice? What does this mean? How much energy does it take to dissociate a CH 4 molecule into atoms?
Homework Exercise 4