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CHEMICAL BONDS

Stability in Bonding. Compounds are substances formed from two or more elements in which the exact combination and proportion of elements is always the same.Example: Copper Sulfate (CuSO4) which makes the Statue of Liberty green in color.When compounds form, the properties of the compound are not

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CHEMICAL BONDS

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    1. CHAPTER 20 CHEMICAL BONDS

    2. Stability in Bonding Compounds are substances formed from two or more elements in which the exact combination and proportion of elements is always the same. Example: Copper Sulfate (CuSO4)– which makes the Statue of Liberty green in color. When compounds form, the properties of the compound are not anything like those of the individual elements. Example: Sodium Chlorine (NaCl) – Sodium is a shiny, soft, silvery metal that reacts violently with water. Chlorine is a poisonous greenish-yellow gas. Table Salt (NaCl) is unreactive and a white crystalline solid.

    3. Stability in Bonding A chemical formula tells what elements a compound contains and the exact number of the atoms of each element in a unit of that compound. Example: Water (H2O) – Contains 2 atoms of hydrogen and 1 atom of oxygen. Notice the subscript 2 written after the H for hydrogen. Subscript means “written below.” A subscript written after a symbol tells how many atoms of that element are in a unit of the compound. If the symbol has no subscript, the unit contains only one atom of that element.

    4. Stability in Bonding In the following compounds, list the number of atoms of each element. NaCl SiO2 C12H22O11 CaO CH3COOH N2O C2H5OH H2SO4 HCl Mg(OH)2 Ca(NO3)2 Al2(SO4)3

    5. Stability in Bonding Why do atoms form compounds? The electric forces between oppositely charged electrons and protons hold atoms and molecules together, and thus are the forces that cause compounds to form. Most of the element in the periodic table can combine with other element to form compounds (Exception = Noble Gases). Noble gases do not form compounds because their outer energy level is full with 8 electrons.

    6. Stability in Bonding To understand the stability of the noble gases, it is helpful to look at electron dot diagrams. Electron dot diagrams show only the electrons in the outer energy level of an atom. Electron dot diagrams contain the chemical symbol for the element surrounded by dots representing its outer electrons.

    7. Stability in Bonding How do you know how many dots to make? Group 1 = 1 Dot Group 2 = 2 Dots Group 13 = 3 Dots Group 14 = 4 Dots Group 15 = 5 Dots Group 16 = 6 Dots Group 17 = 7 Dots Group 18 = 8 Dots

    8. Stability in Bonding Here is the order in which the dots are drawn:

    9. Stability in Bonding Draw electron dot diagrams for the following elements: Sodium Calcium Aluminum Carbon Nitrogen Oxygen Chlorine Neon Tin Potassium Boron Arsenic Sulfur Iodine Krypton

    10. Stability in Bonding An atom is chemically stable when its outer energy level is complete. Has 8 electrons, except Hydrogen and Helium, which will only have 2 electrons. The noble gases are stable because they each have a complete outer energy level. Compounds of these elements rarely form because they are almost always less stable than the original atoms.

    11. Stability in Bonding Atoms with partially stable outer energy levels can lose, gain, or share electrons to obtain a stable outer energy level. They do this by combining with other atoms that also have partially complete outer energy levels. Example: Sodium + Chlorine ? Sodium Chloride

    12. Stability in Bonding When atoms gain, lose, or share electrons, an attraction forms between atoms, pulling them together to form a compound. This attraction is called a chemical bond. Chemical Bond: The force that holds atoms together in a compound.

    13. Types of Bonds When an atom gains or loses an electron it is called an ion. Ion: A charged particles that has either more or fewer electrons than protons. It is the electric forces between the oppositely charged particles (ions) that hold compounds together. Some ions are needed for bodily functions: Example: Iodine – a lack of iodine causes a wide range of problems, most obviously an enlarged thyroid gland (Goiter). Problems can include, mental retardation, neurological disorders, and physical problems.

    14. Types of Bonds

    15. Types of Bonds Remember the following: Metals want to lose electrons. Nonmetals want to gain electrons. How many electrons do metals want to lose? How many electrons do nonmetals want to gain? Group 1 (Alkali Metals) = Lose 1 electron Group 2 (Alkaline Metals) = Lose 2 electrons Group 13 (Boron Group) = Lose 3 electrons Group 14 (Carbon Group) = Lose or gain 4 electrons Group 15 (Nitrogen Group) = Gain 3 electrons Group 16 (Oxygen Group) = Gain 2 electrons Group 17 (Halogens) = Gain 1 electron Group 18 (Noble Gases) = Neither gain nor lose electrons

    16. Types of Bonds When a compound is formed, the elements go from having an unstable outer energy level to a stable outer energy level full of electrons (8). When an atom gains an electron it becomes negatively charged (anion) as it now has more electrons than protons. When an atom loses an electron, it becomes positively charged (cation) as it now has more protons than electrons. The charges of the ions are written after the symbol of the element as a superscript. Superscript means “written above.” Examples: Potassium Ion = K1+, Calcium Ion = Ca2+, Oxygen Ion = O2-, Chlorine Ion = Cl1-.

    17. Types of Bonds

    18. Types of Bonds When atoms attract in this way (opposite charges attract), an ionic bond is formed. Ionic Bond: The force of attraction between the opposite charges of the ions in an ionic compound. In an ionic bond, a transfer of electrons takes place. One element loses electrons while another element gains electrons. When a compound is formed, the opposite charges cancel each other so the compound has no charge (neutral) and is stable. Examples: Sodium Chloride = NaCl, Magnesium Chloride = MgCl2, Calcium Oxide = CaO, Aluminum Sulfide = Al2S3.

    19. Types of Bonds Ionic bonds usually are formed between metals and nonmetals. They also form crystalline solids with very high melting points.

    20. Nonmetals are unlikely to lose electrons. Nonmetals are more stable when they share electrons with each other, rather than by losing or gaining electrons. The attraction formed between atoms when they share electrons is known as a covalent bond. Covalent Bond: Attraction formed between atoms when they share electrons. Covalent bonds form between two nonmetals. A neutral particle that forms as a result of electron sharing is called a molecule. Types of Bonds

    21. Types of Bonds Nonmetals can share one pair of electrons (single bond), two pairs of electrons (double bond), and even three pairs of electrons (triple bond). Examples: Water = H2O; has a two single bonds. Oxygen = O2; has a double bond. Nitrogen = N2; has a triple bond.

    22. Types of Bonds Remember, covalent bonds are formed between nonmetalic elements. Covalent compounds are liquids or gases at room temperature. The electrons in a covalent bond are not always shared equally. The strength of attraction between each atom depends on the size of the atom, the charge of the nucleus, the total number of electrons the atom contains, and how far away from the nucleus the shared electrons are. Example: Hydrogen chloride = HCl; chlorine atoms have a stronger attraction for hydrogens electrons because the chlorine atom is much larger and has more protons. So, the chlorine atom in hydrogen chloride will be slightly negative (?-) and the hydrogen atom will be slightly positive (?+).

    23. Types of Bonds The atom holding the electron more closely always will have a slightly negative charge. The charge is balanced but not equally distributed. This type of molecule is called polar. The term polar means “ having opposite ends.” A polar molecule is one that has a slightly positive end and a slightly negative end. Example: Water = H2O

    24. Types of Bonds Two atoms that are exactly alike can share their electrons equally, forming a nonpolar molecule. A nonpolar molecule is one in which electrons are shared equally in bonds. Such molecules do not have oppositely charged ends. Examples: Oxygen = O2 and Hydrogen = H2.

    25. Writing Formulas and Naming Compounds A binary compound is one that is composed of two elements. Examples: Sodium Chloride = NaCl, Potassium Iodide = KI, Calcium Oxide = CaO. In order to write a formula for a binary compound, you have to know which type (metals or nonmetals) of elements are involved. You also need to know the number of electrons they lose, gain, or share in order to become stable.

    26. Writing Formulas and Naming Compounds The first binary compound we will learn about is a binary ionic compound. A binary ionic compound is composed of a metal and a nonmetal. Remember: Metals want to lose electrons, which makes them positive ions (cations). Nonmetals want to gain electrons, which makes them negative ions (anions). The oxidation number of an element tells you how many electrons an atom has gained, lost, or shared to become stable.

    27. Writing Formulas and Naming Compounds For ionic compounds the oxidation number is the same as the charge of the ion. Examples: Sodium Ion = Na1+, the oxidation number for the sodium ion is 1+. Calcium Ion = Ca2+, the oxidation number for the calcium ion is 2+. Chloride Ion = Cl1-, the oxidation number for the chloride ion is 1-.

    28. Writing Formulas and Naming Compounds The table below shows you the oxidation number of the common elements in binary ionic compounds. The elements with more than one oxidation number, are named using a roman numeral equal to the oxidation number.

    29. Writing Formulas and Naming Compounds Example: Give the symbol and oxidation number for each of the following ions: Potassium Aluminum Copper (II) Copper (III) Phosphorus Fluorine Lead (II) Lead (IV)

    30. Writing Formulas and Naming Compounds In a binary ionic compound, the total numbers of positive charges and negative charges must be equal. All compounds a neutral (have no charge). The formula for a compound can be written given the identities of the compound’s ions.

    31. Writing Formulas and Naming Compounds Three steps to writing chemical formulas for compounds: 1. Write the symbols for the ions side by side. Write the cation first. 2. Cross over the charges by using the absolute value of each ion’s charge as the subscript for the other ion. 3. Check the subscripts and divide them by their largest common factor to give the smallest possible whole number ratio of ions. Then write the formula.

    32. Writing Formulas and Naming Compounds Example: Write formulas for the compounds contains the following elements: Sodium and Chlorine Aluminum and Chlorine Lithium and Nitrogen Lead (IV) and Phosphorus Iron (III) and Oxygen Aluminum and Oxygen Copper (II) and Oxygen Copper (I) and Oxygen

    33. Writing Formulas and Naming Compounds In naming a binary ionic compound, always name it the following way: Cation Name + Anion Name (“-ide”) The ending “-ide” indicates that a compound contains only two elements.

    34. Writing Formulas and Naming Compounds Example: Name the following compounds: Sodium and Chlorine = NaCl Aluminum and Chlorine = AlCl3 Lithium and Nitrogen = Li3N Lead (IV) and Phosphorus = Pb3P4 Iron (III) and Oxygen = Fe2O3 Aluminum and Oxygen = Al2O3 Copper (II) and Oxygen = CuO Copper (I) and Oxygen = Cu2O

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