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SURVEY OF CHEMISTRY I CHEM 1151 CHAPTER 4

SURVEY OF CHEMISTRY I CHEM 1151 CHAPTER 4

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SURVEY OF CHEMISTRY I CHEM 1151 CHAPTER 4

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  1. SURVEY OF CHEMISTRY I CHEM 1151CHAPTER 4 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

  2. CHAPTER 4 FORCES BETWEEN PARTICLES

  3. CHEMICAL BOND - The attractive force that holds atoms together - The result of interactions between electrons in the combining atoms Two types of chemical bonds - Ionic (electrovalent) and covalent bonds

  4. CHEMICAL BOND Ionic Bond - Formed by attraction between two oppositely charged ions - Formed as a result of the transfer of electron(s) from atom(s) to another atom(s) - Often formed between metal and nonmetal ions through electrostatic attraction - Electron transfer

  5. CHEMICAL BOND Covalent Bond - Formed through the sharing of one or more pairs of electrons between two atoms - Always involve two nonmetals - Electron sharing

  6. CHEMICAL BOND Covalent Bond 1s electrons Shared electron pair : H H ∙ ∙ H H Two hydrogen atoms H + H Hydrogen molecule H H

  7. CHEMICAL BOND Covalent Bond - Two neclei attract the same shared electrons to form a covalent bond - Orbitals containing the valence electrons overlap to create a common orbital - The electrons move throughout the common orbital - The electrons are shared by both nuclei

  8. CHEMICAL BOND Two concepts - Valence Electrons - Octet Rule

  9. VALENCE ELECTRONS - Electrons in the outer most shell of an atom - Not all electrons in a given atom participate in bonding - Only valence electrons are available for bonding - For representative and noble-gas elements these electrons are always found in the s or p subshells

  10. VALENCE ELECTRONS - Using electron configuration to determine the number of valence electrons C: 1s22s22p2 O: 1s22s22p4 Na: 1s22s22p63s1 - Using electron-dot structure (Lewis symbol) to designate the number of valence electrons (place first 4 dots separately on four sides and pair up as needed) . .. ∙C∙ :O∙ Na∙ . .

  11. VALENCE ELECTRONS Three important facts about valence electrons - Representative elements in the same group of the periodic table have the same number of valence electrons - The number of valence electrons for representative elements is the same as the group number (with A) in the periodic table - The maximum number of valence electrons for any given element is eight

  12. OCTET RULE - Electrons arranged with 8 valence electrons are more stable than all others - The valence electron configuration of the noble gases are considered the most stable (all have 8 valence electrons; helium has 2) - All noble gases have the outermost s and p subshells completely filled

  13. OCTET RULE - The noble gases are the most unreactive of all elements - Atoms of many elements tend to acquire the 8 valence electron configuration through chemical reactions - Atoms of elements tend to gain, lose, or share electrons to produce a noble-gas electron configuration - This results in the formation of compounds - This tendency is known as the OCTET RULE

  14. CHEMICAL COMPOUND - Consists of two or more different elements chemically bonded together in a fixed proportion by mass Two classes of chemical compounds Ionic and Molecular compounds

  15. CHEMICAL FORMULA - Tells which elements and how many of those elements make up a compound Example The chemical formula of sulfuric acid is H2SO4 Conveys the information that a sulfuric acid molecule contains - 3 different elements: hydrogen (H), sulfur (S), oxygen (O) - 7 atoms: 2 hydrogen atoms, 1 sulfur atom, 4 oxygen atoms - When a particular atom is 1 (as in S above) the subscript is not written

  16. CHEMICAL FORMULA Note the difference between CoCl2 and COCl2 CoCl2 - 2 different elements: cobalt (Co) and chlorine (Cl) - 3 atoms: 1 Co atom and 2 Cl atoms COCl2 - 3 different elements: carbon (C), oxygen (O), and chlorine (Cl) - 4 atoms: 1 C atom, 1 O atom, and 2 Cl atoms Ca(NO3)2 1 Ca atom and 2 NO3- atoms N atoms = 1 x 2 = 2 O atoms = 3 x 2 = 6

  17. IONS - Positively or negatively charged particles - Result from loss (positive) or gain (negative) of electrons - Basic structural units of ionic compounds Two types of ions Cations and Anions

  18. IONS Cation - An ion with a positive charge (loss of electrons) (H+, Na+, Al3+, Ca2+) Anion - An ion with a negative charge (gain of electrons) (O2-, Cl-, Br-, I-) - Charges are represented by superscripts

  19. IONS Generally, metals form cations whereas nonmetals form anions - Chemical properties of ions are different from those of the atoms from which they were derived - Atoms lose or gain electrons to attain the same number of electrons as the noble gas (Group 8A element) closest to it Generally Group 1A elements form 1+ ions Group 2A elements form 2+ ions Group 6A elements form 2- ions Group 7A elements form 1- ions

  20. IONS - Atoms rarely lose or gain more than three electrons - Elements of Group 4A can gain or lose 4 electrons to attain the noble gas configuration - These elements usually form covalent bonds

  21. IONS Monatomic Ion - An ion composed of one atom (Br-, Na+, Mg2+, Al3+, O2-) Polyatomic Ion - An ion (charged particle) composed of two or more atoms covalently bonded together - behave as a single unit (OH-, NH4+, SO42-, CO32-, HCO3-)

  22. IONS Isoelectronic Species - An atom and ion(s) that have the same electron number and configuration Example Ne, F-, O2-, N3-, Na+, Mg2+, Al3+ All these are isoelectronic since each has 10 electrons List 5 ions each that are isoelectronic with Ar, Kr, and Xe

  23. IONIC COMPOUNDS - Ionic bonds are present - Contains both positively and negatively charged ions - Generally composed of metals (positive ions) and nonmetals (negative ions) Examples NaCl, KCl, CaBr2, Na2O

  24. IONIC COMPOUNDS Physical Properties - High melting and boiling points - High hardness - Brittle - Good conductors of electricity when molten - Solid at room temperature and pressure - High solubility in water

  25. IONIC COMPOUNDS - Electron transfer - Metals donate electrons to form positive ions - Nonmetals accept electrons to form negative ions - The electrons lost by the metal are the same ones gained by the nonmetal

  26. IONIC COMPOUNDS - The positive and negative ions attract one another to form ionic compounds - Ions combine in ratios to obtain charge neutrality (net charge = 0) - The symbol for positive ions is always written first

  27. IONIC COMPOUNDS - Ionic compounds do not contain discrete molecules but ordered arrays of positive and negative ions NaCl for example - The formula unit indicates combining ratio (empirical formula) - A given sodium ion has six immediate chloride ion neighbors - A given chloride ion has six immediate sodium ion neighbors

  28. IONIC COMPOUNDS - The charges of ions can be used to depict the empirical formula for ionic compounds - For equal magnitude of charges on cation and anion the subscript on each ion is 1 NaCl: Na+ and Cl- KBr: K+ and Br- MgO: Mg2+ and O2-

  29. IONIC COMPOUNDS - The charges of ions can be used to depict the empirical formula for ionic compounds - For unequal magnitude of charges the charge on one ion is the subscript on the other ion (without the positive or negative sign) AlCl3: Al3+ and Cl- CaCl2: Ca2+ and Cl- Na2S: Na+ and S2- Fe2O3: Fe3+ and O2-

  30. IONIC COMPOUNDS - The charges of ions can be used to depict the empirical formula for ionic compounds - For polyatomic ions each ion is considered as one unit NH4NO3: NH4+ and NO3- (NH4)2CO3 : NH4+ and CO32- Na2CO3: Na+ and CO32- Sr3(PO4)2: Sr2+ and PO43- Ca(OH)2:Ca2+ and OH-

  31. THE MOLECULE - Two or more atoms tightly bound together and behaving as a single unit - Basic structural unit of molecular compounds - The molecule is the limit of physical subdivision (the smallest particle of a compound) - The atom is the limit of chemical subdivision

  32. THE MOLECULE • Homoatomic Molecule • - All atoms present are of the same kind (element) • Examples • H2, O2, N2, Cl2, S8 • Heteroatomic Molecule • - Two or more kinds of atoms are present (compound) • - Two or more elements are present • Examples • H2O, CO2, CH4, HCl

  33. THE MOLECULE Diatomic molecule contains two atoms (HCl, H2, O2) Triatomic molecule contains three atoms (H2O, CO2) Tetratomic molecule contains four atoms (HNO2,SO3) Pentatomic molecule contains five atoms (HNO3, CH4) ETC.

  34. MOLECULAR COMPOUNDS - Compounds composed of molecules - Contain more than one type of atom - Atoms are joined through covalent bonds -Generally composed of nonmetals only Examples H2O, CO2, CH4, NH3

  35. MOLECULAR COMPOUNDS Physical Properties - low melting points - poor conductors of electricity - can be solids, liquids, or gases at room temperature - low solubility in water

  36. BINARY COMPOUNDS - Only two elements are present - Any number of atoms of the two elements may be present Examples NaCl, H2O, CaCl2, NH3, Al2O3

  37. BINARY COMPOUNDS Binary Ionic Compounds - One of the two elements is a metal (cation) and the other is a nonmetal (anion) Examples Al2S3, NaCl, KCl, KBr Binary Molecular Compounds - Both elements are nonmetals Examples H2O, CO2, NO2, SO2, HCl, NH3

  38. LEWIS STRUCTURES - Electron dot structures - Lewis structures involve compounds - Lewis symbols involve individual elements

  39. LEWIS STRUCTURES Ionic Compounds .. .. ∙Cl: [Na]+ Na∙ + [:Cl:]- NaCl .. .. .. ∙Cl: .. .. [:Cl:]- .. CaCl2 ∙Ca∙ + [Ca]2+ .. .. [:Cl:]- .. ∙Cl: ..

  40. LEWIS STRUCTURES Molecular Compounds H∙ ∙H H : H H H H2 .. .. .. .. .. .. F2 :F∙ ∙F: :F : F: :F F: .. .. .. .. .. .. .. .. .. HF H∙ ∙F: H : F: H F: .. .. .. bonding electrons nonbonding electrons

  41. LEWIS STRUCTURES H2O H H H ∙ . .. . H : : OR H O : O O : .. .. .. H ∙ - Oxygen (O) has six valence electrons - Gains two more through electron sharing with H - Achieves a noble-gas configuration

  42. LEWIS STRUCTURES NH3 H ∙ H H . .. . H : : OR H N : N H ∙ N : .. . H H H ∙ - Nitrogen (N) has five valence electrons - Gains three more through electron sharing with H - Achieves a noble-gas configuration

  43. LEWIS STRUCTURES CH4 H ∙ H H H ∙ . .. H : : H OR H C H C ∙ C ∙ .. . H ∙ H H H ∙ - Carbon (C) has four valence electrons - Gains four more through electron sharing with H - Achieves a noble-gas configuration

  44. LEWIS STRUCTURES Bonding Electrons - The pairs of valence electrons involved in the covalent bond formation Nonbonding Electrons (Lone pairs of electrons) - The pairs of valence electrons not involved in electron sharing

  45. SINGLE COVALENT BOND - Two atoms share one pair of valence electrons - Represented by one line

  46. DOUBLE COVALENT BOND - Two atoms share two pairs of valence electrons - Represented by two lines - Approximately twice as strong as a single covalent bond between the same two atoms

  47. DOUBLE COVALENT BOND CO2 - C has four valence electrons and needs four more - Each O atom has six valence electrons and needs two more .. :O::C::O: or O C O .. - Possible for elements that need two electrons to complete their octet

  48. TRIPLE COVALENT BOND - Two atoms share three pairs of valence electrons - Represented by three lines - Approximately thrice as strong as a single covalent bond between the same two atoms

  49. TRIPLE COVALENT BOND N2 - Nitrogen has five valence electrons and needs three more to complete its octet - Each nitrogen must share three of its electrons with the other :N:::N: or :N N: - Possible for elements that need three or more electrons to complete their octet

  50. COORDINATE COVALENT BOND - Both electrons come from only one of the two bonding atoms - Oxygen often forms coordinate covalent bonds : : X + Y X Y filled orbital vacant orbital shared electron pair Hypochlorous acid (HOCl) Chlorous acid (HClO2) .. .. .. .. .. H : O : Cl : H : O : Cl : O : .. .. .. .. .. coordinate covalent bond