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CH. 6

CH. 6. The Structure of Matter. Ch. 6 Section 1 Notes. Compounds and Molecules Pg. 177-182. Chemical Bonds. The forces that hold atoms or ions together in a compound are called chemical bonds . Can be broken, and the atoms rearrange. Chemical Structure.

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CH. 6

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  1. CH. 6 The Structure of Matter

  2. Ch. 6 Section 1 Notes • Compounds and Molecules • Pg. 177-182

  3. Chemical Bonds • The forces that hold atoms or ions together in a compound are called chemical bonds. • Can be broken, and the atoms rearrange

  4. Chemical Structure • The structure of a building is the way the building’s parts fit together • A compound’s chemical structure is the way the atoms are bonded to make the compound

  5. Some models represent bond lengths and angles. • Bond length is the distance between the nuclei of two bonded atoms • If a compound has 3 or more atoms, a bond angle (the angle formed by two bonds to the same atom) tells which way there atoms point.

  6. Atoms are often represented by a ball-and-stick model to help you understand the compounds structure. • Structural formulas also show the structures of compounds. • Chemical symbols are used to represent the atoms • Space-filling model is another way to represent a water molecule. • Shows the space that the oxygen and hydrogen atoms take up, or fill

  7. The chemical structure of a compound determines that properties of that compound. • Compounds with network structures are strong solids. • Ex: quartz • The strong bonds make the melting and boiling point of quartz and other minerals very high

  8. Some networks are made of bonded ions. • Ex: Table Salt (NaCl) • Found in the form of regularly shaped crystals • Made of a repeating network connected by strong bonds • Oppositely attracted ions • High melting and boiling point

  9. Some compounds are made of molecules. • Ex: sugar • Molecules attract each other and form crystals Nitrogen, Oxygen, and Carbon Dioxide --Gases that are made of molecules --Atoms are strongly attracted to each other and are bonded

  10. The strength of attractions between molecules varies. • Sugar, water and Dihydrogen sulfide are all compounds made of molecules but have different properties *The higher the melting point, the stronger the attraction between the atoms

  11. Hydrogen Bond • Oxygen atom of a water molecule is attracted to a hydrogen atom of another molecule • Strong bonds within each water molecule • Weaker attractions between water molecules

  12. Ch. 6 Section 2 Notes • Ionic and Covalent Bonding • Pg. 183-190

  13. Why do Chemical Bonds Form? • Atoms join to form bonds so that each atom has a stable electron configuration. • One similar to a noble gas • There are two kinds of chemical bonding: • Ionic Bonding • Covalent Bonding

  14. Ionic Bonds • Form from the attractions between such oppositely charged ions. • Formed by the transfer of electrons • Oppositely charged ions bond (NaCl)

  15. Ionic compounds are in the form of networks, not molecules. • The formula unit of one sodium ion and one chloride ion is NaCl • NaCl ratio is 1:1 • CaF2 is 1:2 • When melted or dissolved in water, ionic compounds conduct electricity. • Ions are free to move when not is solid form.

  16. Covalent Bonds • Compounds that are made of molecules, such as water and sugar have covalent bonds. • Atoms joined by covalent bonds share electrons. • Usually form between nonmetal atoms.

  17. Can be solid, liquid or gas • Low melting points • MOST do not conduct electricity (not charged) • Example: Cl2 • Each has 7 valence electrons. • Share one electron to have 8 valence electrons and become stable.

  18. Atoms may share more than one pair of electrons. • When drawing the electron dot diagram, a line — means that there are 2 electrons being shared. • If there is two lines ==, that is a double covalent bond (4 electrons being shared) • A triple covalent bond is formed by bonding two nitrogen atoms (total of 6 electrons)

  19. Atoms do not always share electrons equally. • When electrons are shared equally, they are called nonpolar covalent bonds. • Ex: Cl2 • When two atoms of different elements share electrons, the electrons are not shared equally and forms a polar covalent bond. • Ex: NH3

  20. Metallic Bonds • Metals are flexible and conduct electric current well because their atoms and electrons can move freely throughout a metal’s packed structure. • Atoms in metals such as copper form metallic bonds.

  21. Polyatomic Ions • Acts as a single unit in a compound, just as ions that consist of a single atom do. • Groups of covalently bonded atoms that have a positive or negative charge as a group. • Both covalent and ionic bonds • There are many common polyatomic ions.

  22. Parentheses group the atoms of a polyatomic ion. • A polyatomic ions charge applies not only to the last atom in the formula but to the whole ion. • A polyatomic ion acts as a single unit in a compound

  23. Some names of polyatomic anions relate to the oxygen content of the anion. • Most end with –ite or –ate • -ate ending usually used to name an ion that has 3 oxygen atoms • Examples:sulfate (SO42–), nitrate (NO3–), chlorate (ClO3–) • 2 or less oxygen atoms have an –ite ending • Examples:sulfite (SO32–), nitrite (NO2–), chlorite (ClO2–) • Hydroxide and Cyanide are exceptions to the rules.

  24. CH. 6 Section 3 Notes • Compound Names and Formulas • Pg. 191-196

  25. Naming Ionic Compounds • Formed between cations and anions • The names of ionic compounds consist of the names of the ions that make up the compounds.

  26. Names of cations include the elements of which they are composed. • Usually the name of the element • Ex: sodium forms a sodium ion

  27. Names of anions are altered names of elements. • The difference is the name’s ending • Usually with the ending –ide • Compounds with Oxygen atoms have –ate, or –ite endings • An ionic compound must have a total charge of zero.

  28. Some cation names must show their charge. • Transition metals may form several cations (each will have a different charge). • Iron forms a +2 ion AND a +3 ion • This is shown by placing the charge of the cation as a Roman numeral in parentheses. • Iron (II) ion and Iron (III) ion • FeO --- Iron (II) Oxide • Fe2O3 --- Iron (III) Oxide

  29. Determining the charge of a transition metal cation. • The total charge of the compound MUST be zero. • Fe2O3 • Three oxide ions have a total charge of 6-. (each oxygen ion has a charge of 2- 2-(3)=6-) • So, the total charge of the cation must be 6+

  30. Writing Formulas for Ionic Compounds • If you are given the compound’s name: you can find the formula • If you are given the formula: you can find the charge of each ion

  31. Naming Ionic Compounds Rules • Calcium Chloride • Ca, Cl • Ca+2 , Cl -1 • CaCl2 • This is not a polyatomic Ion • This is a polyatomic Ion • If you are given the Name: • 1. Find the symbol of each element • 2. Find the charge of each ion • 3. Criss-cross Method • 4. If one of the ions is a Polyatomic Ion, put parentheses around it!!

  32. Naming Ionic Compound Rules: • AgF • Since there is no subscript number the charges for both must be 1. • Ag is Silver, F is Flourine. • F is in group 17 and has a -1 charge so, Ag is the cation. • Silver Flouride • If you are given the formula: • 1. Determine if the FIRST ion is a Transition metal. If so, you MUST find it’s charge! • 2. Find the name of each of the ions • 3. The cation is the same as it is on the periodic table • 4. The anion has an –ide ending (unless it is a polyatomic ion)

  33. CrO2 1:2 ratio Cr O2 Cr +4 O-2 Cr2O4 Ratio is 2:4 Reduce the ratio to 1:2 • To find the charge of ions in a chemical formula: • Determine the ratio of the given formula • Separate the ions • Determine each of their charges • If the cation is a transition metal, use the criss-cross method and then look at it’s ratio. • Compare to the original ratio. What ever you do to the first element, you must do the the 2nd.

  34. Math Skills “Writing Ionic Formulas” Practice Problems 1-3 Pg. 193 • Lithium oxide • Li+1 O-2 • Li2O • Beryllium chloride • Be+2 Cl-1 • BeCl2 • Titanium (III) nitride • Ti+3 N-3 • TiN

  35. Naming Covalent Compounds • For covalent compounds of two elements, numerical prefixes tell how many atoms of each element are in the molecule. • Numerical prefixes are used to name covalent compounds of two elements. • If there is only one atom of the first element, the name does not get a prefix.

  36. BF3 • Boron Trifluoride • N2O4 • Dinitrogen tetroxide

  37. Empirical Formulas • Chemical formulas that are unknown are determined by figuring out the mass of each element in the compound. • Once the mass of each element is known, scientists can calculate the compound’s empirical formula, or simplest formula. • An empirical formula tells us the smallest whole-number ratio of atoms that are in a compound.

  38. Different compounds can have the same empirical formula. • Molecular formulas are determined from empirical formulas. • A compound’s molecular formula tells you how many atoms are in one molecule of the compound. • Masses can be used to determine the empirical formula. • Convert the masses to moles. Then, find the molar ratio to give you the empirical formula.

  39. Pg.196 • Math Skills “Finding Empirical Formulas” • One mole of an unknown compound has 36.04 g of Carbon and 6.04g of hydrogen. What is the compound empirical formula.

  40. Section 3 Review # 1, 5 • 5. Determine the chemical formulas for the following ionic compounds. • Magnesium sulfate • MgSO4 • Rubidium bromide • RbBr • Chromium(II) fluoride • CrF2 • Nickel(I) carbonate • Ni2CO3 • Name the following ionic compounds, and specify the charge of any transition metal cations. • FeI2 • Iron(II) Fluoride • MnF3 • Manganese(III)Flouride • CrCl2 • Chormium(II) Chloride • CuS • Copper(II) Sulfide

  41. Ch. 6 Section 4 Notes • Organic and Biochemical Compounds • Pg. 197-204

  42. Organic Compounds • An organic compound is a covalently bonded compound that contains carbon. • Most contain hydrogen. • Oxygen, nitrogen, sulfur, and phosphorus can also be found in organic compounds.

  43. Carbon atoms form four covalent bonds in organic compounds. • A compound made of only hydrogen and carbon atoms is known as a hydrocarbon. • Methane, CH4 is an example • There are four single C-H bonds • A carbon atom may never form more than 4 bonds at a time.

  44. Alkanes are hydrocarbons that have only single covalent bonds. • Can have C-C bonds as well as C-H bonds • Methane is the simplest alkane

  45. Arrangements of carbon atoms in alkanes. • The carbon atoms in methane, ethane, and propane are all bonded in a single line because that is their only possible arrangement. • If there are more than 3 bonded carbon atoms in a molecule, the carbon atoms do not have to be in a single line.

  46. IF they are in a single line: the alkane is a normal alkane, or n-alkane. • The condensed structural formula shows how the atoms bond.

  47. Alkane chemical formulas usually follow a pattern. • Except for cyclic alkanes • The # of Hydrogen atoms is always 2 more than 2x the # of carbon atoms • CnH2n+2

  48. Alkenes have double carbon-carbon bonds. • Hydrocarbons • Have at least one double covalent bond between carbon atoms. C=C • Replace the –ane ending with –ene. • Simplest alkene is ethene (ethylene)

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