Valence Electrons and Chemical Bonds

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1 Valence Electrons and Chemical Bonds A chemical bond is the force that holds two atoms together. Chemical bonds form by the attraction between the positive nucleus of one atom and the negative electrons of another atom. Atom s try to form the octet the stable arrangement of eight valence electrons in the outer energy level by gaining or losing valence electrons. This is the Octet Rule Ion Formation

2 Why do atoms bond? The stability of an atom, ion or compound is related to its energy: lower energy states are more stable. Metals and nonmetals gain stability by transferring electrons (gaining or losing) to form ions that have stable noble-gas electron configurations. Another way atoms can gain stability is by sharing valence electrons with other atoms. The Covalent Bond

3 Lewis Dot Diagrams or Electron Dot Formulas Diagrams that show valence electrons in the atoms of an element as dots around the symbol of the element. H Mg Ga C As O Br Ionic Bonds and Ionic Compounds

4 Metallic Bonding Metallic Bond holds the atoms of metals together When bonding Metals form crystal lattices known as cations surrounded by a sea of freely moving valence electrons. A metallic bond is formed between all metals. Examples include a piece of Copper, Zinc, Sodium, Iron. Any metal Metallic bonds result from the attraction between metal atoms and the surrounding sea of delocalized electrons Valence electrons can move freely around the whole metal structure they are not confined to any one atom Ionic Bonds and Ionic Compounds

5 Metallic Bonds Metallic Bonds and the Properties of Metals

6 This model explains many properties of Metals The mobile electrons can enter/leave the metal structure, so metals are good conductors of heat and electricity! Metal cations are not locked into any crystal structure, so they can slide past each other when stressed. That makes metal malleable and ductile. The de-excitation ( The electron falling back down to a lower energy level) is responsible for the shiny (luster) appearance of metals. Ionic Bonds and Ionic Compounds

7 Metallic Bonds Boiling points are much higher than melting points because of the energy required to separate atoms from the groups of cations and electrons. Metallic Bonds and the Properties of Metals

8 Metallic Bonds Metals are malleable because they can be hammered into sheets. Metals are ductile because they can be drawn into wires. Mobile electrons surrounding positively charged nuclei make metals good conductors of electricity and heat. As the number of delocalized electrons increases, so does hardness and strength. Metallic Bonds and the Properties of Metals

9 Ionic Bonding The Covalent Bond

10 Formation of an Ionic Bond Ionic bonds form when atoms transfer electrons. The electrostatic force that holds oppositely charged particles together in an ionic compound is called an ionic bond. Compounds that contain ionic bonds are called ionic compounds. Ionic Bonds and Ionic Compounds

11 Properties of Ionic Compounds Positive and negative ions exist in a ratio determined by the number of electrons transferred from the metal atom to the non-metal atom. The repeating pattern of particle packing in an ionic compound is called an ionic crystal. The strong attractions among the positive and negative ions result in the formation of the crystal lattice. A crystal lattice is the three-dimensional geometric arrangement of particles, and is responsible for the structure of many minerals. Ionic Bonds and Ionic Compounds

12 Physical Properties of Ionic Compounds Strong attraction between ions Soluble in water Conduct electricity in solution Conduct electricity when molten High melting points High boiling points Hard but brittle Solid at room temperature Ionic Bonds and Ionic Compounds

13 Physical Properties of Ionic Compounds In a solid, ions are locked into position and electrons cannot flow freely solid ions are poor conductors of electricity. Liquid ions or ions in aqueous solution have electrons that are free to move, so they conduct electricity easily. An ion in aqueous solution that conducts electricity is an electrolyte. This figure demonstrates how and why crystals break when an external force is applied. Ionic Bonds and Ionic Compounds

14 How are electron dot formulas for ionic bonds constructed? Electron dot formulas show the valence electrons and charges of these ions and may be used to illustrate the ionic bonds.

15 Examples

16 Covalent Bonding In a Covalent Bond the valence electrons are shared between non-metal atoms The Covalent Bond

17 Properties of Covalent Bonding Two non-metals bonded together Relatively weak bonds Usually a gas or liquid at room temperature Does not conduct electricity in solution Low melting point Low boiling point Soluble in alcohol and insoluble in water The Covalent Bond

18 Types of Covalent Bonds Nonpolar Covalent bond- Covalent bond in which the bonding electrons are shared equally by the bounded atoms, resulting in a balanced distribution of an electrical charge. Polar Covalent bond- Covalent bond in which the bonded atoms have an unequal attraction for shared electrons. Each atom has a charge Partial positive- (+δ) Partial negative- (- δ) The Covalent Bond

19 Covalent Bonding The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together. The Covalent Bond

20 Covalent Bonding If the difference in electronegativities is between: 1.8 to 4.0: Ionic 0.4 to 1.7: Polar Covalent 0.0 to 0.3: Non-Polar Covalent Example: NaCl Na = 0.9, Cl = 3.0 Difference is 2.1, so this is an ionic bond! Example: H 2 O H = 2.1, O = 3.5 Difference is 1.4, so this is a polar covalent bond! The Covalent Bond

21 Diatomic Molecules Diatomic Molecules are nonpolar covalent compounds. A diatomic molecule is a molecule only containing two atoms The seven diatomic molecules are H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 These diatomic molecules are never by themselves.

22 Single Covalent Bond Single Covalent Bonds -A single covalent bond results when two atoms share one pair of electrons, as in the case of hydrogen gas, which is a diatomic molecule. The Covalent Bond

23 Examples: F 2, HBr, Cl 4 (contd.)

24 Double Covalent Bonds- When atoms bond by sharing two pairs of electrons, the result is a double covalent bond, as in a molecule of carbon dioxide, CO 2. The double bond is shown by four dots or two dashes.

25 Triple Covalent Bonds- When atoms bond by sharing three pairs of electrons, the result is a triple covalent bond, as in a molecule of nitrogen gas, N 2. The triple bond is shown by six dots or three dashes.

26 Violations of the Octet Rule Usually occurs with B and elements of higher periods. Common exceptions are: Be, B, P, S, and Xe. Be: 4 B: 6 P: 8 OR 10 SF 4 BF 3 S: 8, 10, OR 12 Xe: 8, 10, OR 12

27 Lewis Structures Single Bond HF H 2 O

28 Lewis Structures Double & Triple Bond C 2 H 4 N 2

29 Bond Polarity H 2 O is POLAR because it has a positive end and a negative end. (difference in electronegativity) O has a greater share in bonding electrons than does H. O has slight negative charge (-δ) and H has slight positive charge (+ δ)

30 Bond Polarity This is why oil and water will not mix! Oil is nonpolar, and water is polar. The two will repel each other, and so you can not dissolve one in the other

31 Bond Polarity Like Dissolves Like Polar dissolves Polar Nonpolar dissolves Nonpolar Polar also dissolved Ionic Ex: water and salt

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