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Chapter 1 Chemical Bonding

Chapter 1 Chemical Bonding. All matter is made up of atoms . Atoms are the basic building blocks of all the substances in the universe. Elements are substances made of only one kind of atom. There are 109 naturally occurring elements in the universe. .

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Chapter 1 Chemical Bonding

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  1. Chapter 1 Chemical Bonding

  2. All matter is made up of atoms. Atoms are the basic building blocks of all the substances in the universe.

  3. Elements are substances made of only one kind of atom. There are 109 naturally occurring elements in the universe.

  4. Elements can combine with other elements to produce new substances called compounds. Compounds contain two or more kinds of atoms that are chemically joined together. The combining of elements to form new substances is called chemical bonding.

  5. Structure of an atom Each atom contains a positively charged center called a nucleus. The nucleus contains two types of particles: a. protons- positive charge, + b. neutrons- neutral charge, 0

  6. Outside of the nucleus are negatively charged particles called electrons, which are located in an “electron cloud”. The electron cloud contains many different energy levels that can hold only a certain amount of electrons.

  7. The first energy level can only hold 2 electrons The second energy level can hold 8 electrons The third energy level can hold 8 electrons

  8. Electrons in the outermost energy level are called valenceelectrons. Electrons in this energy level are the furthest away from the nucleus, and are involved in bonding. Atoms that have their outermost energy level filled are very stable. They usually do not form compounds or form chemical bonds.

  9. On pages 130-131 of your book is the Periodic Table of theElements. All 109 elements are listed on this table in a specific way. Vertical columns are called families (18), horizontal columns are called periods (7).

  10. The elements in family 18 (He, Ne, Ar, Kr, Xe, Rn) do NOT form chemical bonds. Their outermost energy levels are filled with electrons. These are stable elements!

  11. Electrons and Bonding As we learned before, the number of electrons in the outermost energy level determines whether or not the atom will form chemical bonds. Atoms that do not have complete valence energy levels will undergo chemical bonds to fill these energy levels and become stable. Example: Fluorine (F) has 7 valence electrons. It will gain one electron to fill the outermost energy level (8).

  12. Ionic Bonds Bonding that involves the transfer of electrons between 2 atoms is called ionicbonding. Because this type of bonding involves the transfer of electrons, one atom gains an electron and the other atom loses an electron.

  13. REMEMBER! Because electrons are negatively charged, when an atom gains an electron, it receives a negative charge (-1). When an atom loses an electron, it loses a negative charge (+1) Example: Fluorine has 7 valence electrons. It will gain an additional electron to fill this outer level and become a fluoride ion (F-1). An ion is any atom that has a + or – charge.

  14. In nature, opposite charges attract one another. The sodium ion has a positive charge (Na +1) while the fluoride ion has a negative charge (F -1). The attraction between the + and – ions will result in an ionic bond so that the charges are balanced. The compound NaF (sodium fluoride) will be formed. Na+1 + F-1 NaF

  15. The arrangement of ions in an ionic compound appears in a regular, repeating pattern called a crystal lattice. This pattern results in the formation of crystals.

  16. Covalent Bonds Bonding in which electrons are shared rather than transferred is called covalentbonding. By sharing electrons, each atom fills up its outermost energy level and becomes stable!

  17. Scientists represent the electron sharing that takes place in a covalent bond using an electron-dot diagram. In this system, we use the chemical symbol for the element and place dots around the symbol to represent the number of valence electrons.

  18. Elements such as H, Cl, O, and N often form covalent bonds with other atoms of the same element. This would give us H2, Cl2, O2, and N2. Elements that can bond in this way are called diatomicelements.

  19. The combination of atoms formed by a covalent bond is called a molecule. A molecule is made up of 2 or more types of atoms that are covalently bonded together. Molecules are represented by chemical formulas: H20 (2 atoms of H, 1 atom of O) CO2 (1 atom of C, 2 atoms of O)

  20. Metallic Bonds Metals are elements that give up electrons easily. When 2 or more atoms of a metal combine, they form a metallic bond. In this type of bonding, the valence electrons of several metal atoms are shared as a common electron cloud.

  21. The electrons involved in a metallic bond are often described as a “sea of electrons” because they are able to freely move from atom to atom, much like water moves.

  22. Properties of Metals: • malleable- ability to be hammered • into thin sheets without breaking. • 2. ductile- can be drawn into thin wire • 3. Flexibility • 4. good conductors of heat and electricity

  23. Predicting Types of Bonds The placement of the elements on the periodic table indicated what types of bond it will usually form. Elements on the left and in the center of the table are metals. Elements on the right hand side of the table are nonmetals.

  24. A. Metal + Metal = Metallic Bond B. Metal + Nonmetal = Ionic bonds. C. Nonmetal + Nonmetal = Covalent Bond.

  25. We have already learned that the valence electrons of an atom determines whether or not the atom will form a bond. We know that an atom will gain, lose, or share these electrons when it forms a compound. The number of electrons gained, lost, or shared when an atom forms a chemical bond is called its oxidation number.

  26. Example: Na has 1 valence electron. When it forms a bond, Na will lose this electron to become +1. So, the oxidation number of Na is +1. Example: Cl has 7 valence electrons. When it bonds it will gain an electron and have a -1 charge. Therefore, the oxidation number of Cl is -1.

  27. Oxidation numbers can be used to predict how atoms combine and what the formula will be. RULE: the sum of the oxidation numbers for a compound must be zero!!!

  28. Example problems Mg has an oxidation number of +2. Cl has an oxidation number of -1. What would be the formula for magnesium chloride? We would need 2 Cl molecules to equal the +2 oxidation number of Mg.  MgCl2

  29. Na has 1 valence electron, Oxygen has 6 valence electrons. What would be the oxidation numbers for Na and O? What is the formula for sodium oxide? Na: +1 oxidation number O: -2 oxidation number Formula: Na2O

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