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Unit 6 Bonding

Unit 6 Bonding. How elements interact. The Octet Rule. Atoms bond when their valence electrons interact. Atoms bond to obtain a complete outer shell of 8 electrons. (Hydrogen and helium only need 2 outer electrons.) In this configuration they are stable.

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Unit 6 Bonding

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  1. Unit 6 Bonding How elements interact.

  2. The Octet Rule • Atoms bond when their valence electrons interact. • Atoms bond to obtain a complete outer shell of 8 electrons. (Hydrogen and helium only need 2 outer electrons.) In this configuration they are stable. • Energy is stored in chemical bonds. When you break a bond, energy is absorbed. When you form a bond, energy is released.

  3. Ionic Bonds • Formed between metals and nonmetals. • Ionic bonds are formed by the transfer of electrons. • The resulting negative and positive ions are held together by electrical attraction. • Atoms on opposite sides of the periodic table are likely to form ionic bonds. Metals lose, nonmetals gain. • Examples:NaCl CaF2

  4. Metallic Bonds • Metals, like copper or gold, form metallic bonds between their atoms. • In metallic bonds the electrons are free to move from atom to atom. • In other words a sea of electrons surrounds the atom. • This explains why metals conduct electricity. • This also explains why metals are ductile and malleable. The atoms can slide past each other without breaking bonds.

  5. Covalent Bonds • Covalent bonds are formed by the sharing of electrons. • Covalent bonds are formed between non-metal atoms. • A double bond is formed if 2 pairs of electrons are shared. Examples of compounds with double bond are: O2 and CO2 • A triple bond is formed if 3 pairs of electrons are shared. Examples of compounds with triple bonds are: N2.

  6. Covalent Bonds con. • Atoms do not share electrons equally. When one atom attracts the electrons more than the other the bond is called polar covalent. • Elements in the upper right hand corner of the periodic table, excluding the noble gases, usually have the greater attraction for electrons. Fluorine will attract electrons more than any other element.

  7. Covalent Bonding Below is a Lewis dot diagram for Cl bonded to Cl to form a diatomic compound Cl2. Cl atoms are sharing the one pair of electrons.

  8. Hydrogen Bonding Hydrogen bonding is an example of a strong intermolecular force. It is the strongest of all the bonds. Why do you think this would be a strong bond? (Hint: think about electronegativity differences.) What are the three molecules that bond with hydrogen to make hydrogen bonds?

  9. Compounds • A compound consists of two or more elements chemically bonded together. • Differences between compounds and mixtures. CompoundsMixtures bonded not bonded definite proportions varying proportions 2 or more elements contains elements/cmpds definite properties properties vary w/ compo chemical formula no chem formula

  10. Chemical Compounds • Chemical formulas can be written for compounds. The chemical formulas give the elements and their proportions: • Ex. H2O Contains 2 H atoms for every 1 O atom • A molecule is the smallest unit of a compound that still has the properties of that compound. • Chemical formulas show the molecule’s structure or arrangement of atoms. This is very important for biological and pharmaceutical companies.

  11. Structure and Properties Ionic Solids: • Ions held together by the strong attraction of opposite charges. • Occur when metals are bonded to nonmetals. • Have a crystalline structure • High melting and boiling points • Poor conductors of electricity in the solid state because the charged ions are rigid in place.

  12. Ionic Solids con. • In aqueous solutions the ions separate and are free to move. This makes their aqueous solutions good conductors of electricity. • They will also conduct electricity when melted. • Examples: NaCl NaF CaF2

  13. Network Solids • A large network of atoms covalently bonded together in a fixed pattern. • The bonds are very strong • The resulting solid is strong and hard • They are poor conductors • High melting and boiling points b/c the bonds are hard to break. • ONLY 2 Examples: diamond = carbon and SiO2 = quartz

  14. Molecular Substances • Made up of molecules • Can be gases, solid or liquid • Relatively low melting and boiling point • Poor conductors • Examples: H2O, CH4, sugar C11H22O11

  15. Polyatomic Ions – TABLE E!! • Some compounds have both ionic and covalent bonds. Ex: NaOH • The covalently bonded ions form a unit called a polyatomic ion. • A polyatomic ion contains two or more covalently bonded atoms • All the atoms share the charge of a polyatomic ion. • A polyatomic ion may bond with another polyatomic ion or a single “atom” ion. The resulting bond will be ionic.

  16. Compound Names and Formulas • Naming ionic compounds • The name of the metal is listed first • The name of the nonmetal follows with its ending changed to –ide. • Indicate the oxidation state of transition metals by using roman numerals. Ex: NaCl Sodium chloride

  17. Naming Polyatomic Ionic Compounds 4 Steps to Naming Polyatomic Ionic Compounds Step 1: Write the name of the cation ion first Step 2: Write the name of the polyatomic ion second Step 3: Put them together to form a name Step 4: Indicate the oxidation state of transition metals by using roman numerals Example: Na2CO3

  18. Naming Covalent Molecules Use the following steps: Step 1: Find the prefixes that correspond to the number of atoms present in the anion. • Mono – 1, di-2, tri-3, tetra – 4, penta – 5, hexa – 6, hepta – 7 , octa - 8 Step 2: Name the molecule with the cation first and add “ide” to the anion with the prefix

  19. Converting an Ionic Compound Name to a Formula • List the symbols for each ion with the correct charge. • Write the symbols with the metal ion first. • “Criss-cross” the charges to determine the number of each ion that is required to make the compound neutral. • Check to be sure that the compound is neutral. Ex: Aluminum fluoride Al+3F-1  AlF3

  20. Polarity in Covalent Molecules Electrons are not always shared ____________ ______________________ will determine where the shared electrons spend ________ of their time High electronegativity __________ the electrons _____________ that atom making the molecule ___________

  21. Polar vs. Nonpolar • Polar Molecules: • Electron pairs are shared _______________ • Molecules do not have _______________ • Delta symbol helps show _______________ • Nonpolar Molecules: • Electron pairs are shared _______________ • Molecules have ____________

  22. Determining Polarity • Molecules are _______ when an atoms have ___________ in electronegativities greater than _____ • Electrons __________ towards the atom with higher ________________________ • Ex: Water Oxygen is more electronegative than Hydrogen so the electrons stay mostly by O

  23. Nonpolar Molecules • Electronegativities are the _________ on each end or the ___________ is less than _________ • Oxygen atoms have an ___________ pull on the electrons causing ___________sharing between all atoms • Symmetry is __________

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