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Chemical Bonding

Chemical Bonding. Unit IV. I. Chemical Bonds :. are attractive forces that hold atoms and/or compounds together. result from the simultaneous attraction of an atom’s positively charged nucleus for other atoms negatively charged electrons.

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Chemical Bonding

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  1. Chemical Bonding Unit IV

  2. I.Chemical Bonds: • are attractive forces that hold atoms and/or compounds together. • result from the simultaneous attraction of an atom’s positively charged nucleus for other atoms negatively charged electrons. • result when electrons are transferred (ionic) and shared (covalent) between atoms. • results in the increase in the chemical stability of atoms when energy is released (exothermic).

  3. II. Chemical Bonds and Energy • a) To break a chemical bond, the given chemical must absorb energy from the environment. • BOND BREAKING IS AN ENDOTHERMIC PROCESS. • Absorb energy • b) When chemical bonds are formed, the stability of the reactants is generally increased. The reactants stability is increased as they release energy. • BOND MAKING IS AN EXOTHERMIC PROCESS. • Release energy NOTE: WHEN CHEMICALS CONTAIN LARGE AMOUNTS OF ENERGY, THEY ARE CONSIDERED TO BE UNSTABLE and are VERY REACTIVE. WHEN CHEMICALS CONTAIN SMALL AMOUNTS OF ENERGY, THEY ARE CONSIDERED TO BE STABLE and are UNREACTIVE.

  4. When atoms release much energy as a result of bond formation, the bond is considered to be strong and stable. Thus in order to break the bond, the substance must absorb much energy. • When atoms release little energy as a result of bond formation, the bond is considered to be weak and unstable. Thus in order to break the bond, the substance must absorb little energy.

  5. Aim: What is the difference between ionic and covalent bonds? Do Now: Take out your labs and read the procedure Homework: Covalent naming sheet

  6. IV. Types of Bonds • Intramolecular Forces – bonds between atoms • a) Ionic Bonds: Characteristics • 1. Ionic bonds exist between ions (charged particles). 2. Ionic bonds involve the transfer of electrons (one atom gains and the other loses). • 3. Ionic bonds are generally formed between a metal (electron donor) and a nonmetal (electron acceptor). • 4. Ionic bonds are considered to be very strong bonds. 5. Ionic bonds are predicted by a difference in electronegativity greater than 1.7.

  7. b) Ionic Compounds: Characteristics • 1. High melting and boiling points. • 2. Solids at STP (standard temperature and pressure). • 3. Form crystal lattice structures. • 4. Conduct electricity in the liquid and aqueous phases only. • 5. Have regular geometric arrangements. • Examples include: • NaCl (Sodium Chloride) • LiCl (Lithium Chloride) • KF (Potassium Fluoride) • MgBr2 (Magnesium Bromide) Metal Nonmetal

  8. Representation of an Ionic Bonding [Na]+ [Cl]- [Cl]- [Mg]2+ [Cl]-

  9. Aim: what is the difference between covalent and ionic bonds? Do Now: what elements want electrons and why? Homework:

  10. c) Covalent Bonds: Characteristics 1. Are bonds between nonmetals. 2. Involve the equal and unequal sharing of electrons. 3. Are considered to be relatively weak bonds. * *Exceptions are network solids (diamonds, graphite, silicon dioxide, asbestos, silicon carbide)

  11. Aim: what is the difference between polar and non-polar covalent bond? Do Now: what causes electrons to be pulled closer to one element rather then another Homework: Castle Learning assignment #4

  12. 4.TYPES: • polar covalent bonds – unequal sharing of electrons; occurs between two different nonmetals; electronegativity difference is between 0 and 1.7. • These compounds are not symmetrical ex: H2O, SO2, CH4 Sulfur dioxide molecule

  13. nonpolarcovalent bonds - equal sharing of electrons; occurs between two of the same nonmetals; electronegativity difference is equal to 0. • Symmetrical Compounds • ex: H2, O2, N2, Cl2, Br2, I2, F2 - DIATOMS • -BrINClHOF • coordinate covalent bonds – sharing of electrons whereas both of the electrons of the shared pair are donated by the same atom (not required for the regents exam).

  14. Note: • Most atoms will share a single pair of electrons that is represented by a single bond (a dash).

  15. Other atoms have the ability to share two or three pairs of electrons representing a double bond (two dashes) or a triple bond (three dashes). . . . . : : O O Double bond (4 electrons shared) : : N N Triple bond (6 electrons shared)

  16. 5. Characteristics of Covalent Compounds • Are also known as molecular substances. • Exist in all three phases at STP. • Poor conductors of heat and electricity (good insulators). • Have low melting and boiling points.

  17. NOTE: Network Solids are covalent compounds that are extremely hard and have very high melting and boiling points. These represent exceptions to the general rules of covalent compounds. Examples include: diamonds, graphite, SiO2, and SiC. Graphite Diamond Network Solids (covalent compounds)

  18. Aim: What is a metallic bond? Do Now: List the difference between covalent and ionic bonds? Homework: Text Book Pg 381 Questions 1-4

  19. d) Metallic Bonds - very strong bonds between metal ions. • “positive ions immersed in a sea of mobile electrons”. • Bonds between metals. • ex: Ag(s), Mg(s), Ca(s) • Metallic Compounds are lustrous, ductile, malleable, and excellent conductors of heat and electricity. • Solids at STP (except Mercury (Hg) which is a liquid).

  20. V. The Octet Rule • Electrons are found outside of the nucleus of atoms. • Electrons are arranged around the atom’s nucleus according to the amount of energy they possess. • Electrons are found in energy levels. • The outermost energy level (valence shell) for a given atom contains valence electrons. • In the case of most atoms of the periodic table, for the atom to be stable, it must contain a total of 8 electrons in its valence level. • To become stable atoms will gain, lose, or share electrons to obtain this valence shell configuration.

  21. More precisely, for an atom to be stable, it must have an electron configuration similar to that of a noble gas. • Exceptions to the Octet Rule include: hydrogen, helium, lithium, beryllium, sulfur, nitrogen.

  22. Aim: What is a Dipole Dipole bond? Do Now: How can a molecule be non-polar but have polar bonds? Homework: Text Pg 394 1-8

  23. Aim: What is a hydrogen bond? Do Now: What creates poles on a molecule? Homework: Castle Learning #5 and Quiz Friday

  24. VI. Bonds Between Molecules • Intermolecular Forces that are weaker in comparison to intramolecular forces (ionic, covalent, and metallic bonds). a) Dipole-Dipole Interactions • Covalent compounds that have unequal sharing of electrons (polar covalent bonds) are considered to be polar molecules. • The unequal sharing of electrons results in molecules that have positively charged and negatively charged regions. • The positive region of one polar molecule will be attracted to the negative region of another polar molecule (vice versa).

  25. From: http://www.geo.arizona.edu/xtal/geos306/9_7.jpg

  26. Hydrogen bonds between water molecules

  27. Hydrogen bonds. An example of dipole-dipole interactions. • Oxygen has a greater electronegativity than hydrogen. Thus, the electrons are shared unequally. • The shared electrons are drawn closer to the oxygen atoms giving the them a slight negative charge. The hydrogen atoms have a slight positive charge. • The negatively charged oxygen of one water molecule is attracted to the positively charged hydrogen of the other water molecule.

  28. 1. Molecular Polarity and Symmetry Polar and Nonpolar Molecules: • A polar molecule is one that has an uneven electron distribution resulting in a net positive and negative charge in different parts of the compound. • A non-polar molecule is one that has an even distribution of electrons resulting in no net charge. • If a compound has non-polar bonds (as with the diatomic molecules), the molecule is considered to be non-polar.

  29. In most cases, if a molecule contains a polar covalent bond, it is considered to be a polar molecule. • The exception exists when the molecule is completely symmetrical. • A symmetrical molecule will have an even distribution electrons resulting in no net charge. Carbon Dioxide (CO2): Linear Structure Methane (CH4): Tetrahedral Structure

  30. The four molecules that have polar covalent bonds but are non-polar molecules due to their symmetry are: CH4 - Methane CCl4 - Carbon Tetrachloride CF4 - Carbon Tetrafluoride CO2 - Carbon Dioxide

  31. Aim: what is the difference between empirical and molecular formulas? Do Now: how do water molecules line up around dissolved salt in waster NaCl? Draw this. Homework: Text Book

  32. b) Van der Waals Forces • Weak forces of attraction between nonpolar compounds. • Remember, most nonpolar compounds contain nonpolar bonds which equally distribute electrons between the atoms. • Forces of attraction between: • Monoatomic Molecules (Noble Gases) • Diatomic Molecules • Other select nonpolar substances.

  33. c) Molecule-Ion Attractions • Forces of attraction between ions and polar covalent compounds. Example: Which of the following contain molecule-ion attractions? 1. CaCl2 (s) 2. CO2 (g) 3. NaCl (aq) 4. Ag (s) Answer: Choice 3. Sodium chloride (being a salt, ionic compound, and an electrolyte) will dissolve, dissociate, and form ions in water. Water is a polar covalent molecule. Thus…. Molecule-Ion Attraction. Hint: Always look for the choice that has an ionic compound in the aqueous state.

  34. Types of Bonds in Compounds with Polyatomic Ions • Compounds with two different element are called binary compounds. The types of bonds found between atoms in a binary compound are dependent on the types of elements. • Examples of such compounds include: CO2 NaCl CaF2 Li2O Carbon Dioxide Lithium Oxide Sodium Chloride Calcium Fluoride • Compounds with three or more different elements are called ternary compounds. • If a ternary compound has a metal and a polyatomic ion (see reference table E), it will have both ionic and covalent bonds.

  35. Na2SO4 KNO3 Potassium Nitrate Sodium Sulfate Li3PO4 Mg(OH)2 Magnesium Hydroxide Lithium Phosphate

  36. Ionic Bond (bond between a metal and a non-metal) Covalent Bond (bond between two non-metals)

  37. VII. Formulas • Chemical Formula: represents the type and number of element(s) in a chemical compound. • The type of element is represented by its chemical symbol. • The number of atoms of the element in the compound is represented by its subscript. NOTE: Subscripts are the numbers found to the lower right of a given chemical symbol. If a subscript is not present assume that it is a number one. Li3PO4 3 atoms 1 atoms 4 atoms Lithium = Li Phosphorus = P Oxygen = O

  38. Types of Formulas • Molecular Formulas • Represents the type and number of elements in a covalent compound. • The subscripts do not need to be reduced in a molecular formula. Ex:C2H4Ethene C6H12O6 Glucose

  39. 2. Empirical Formulas • Represents the lowest whole number ratio of atoms in a compound. • For ionic compounds, the formula is always empirical. This is due to the fact that one must reduce the subscripts in an ionic compound to the lowest whole number ratio. Ex: Calcium Sulfide Ca S Ca+2 S-2 Crisscross Method Ca2 S2 Disregard the signs. Ca S Reduce

  40. For covalent compounds (molecules), reducing its molecular formula to an empirical formula only shows the ratio of elements within the compound. Ex: C6H12O6Molecular Formula C1H2O1Empirical Formula Ratio: C:H:O = 1:2:1

  41. H2O Molecular and Empirical Formula

  42. Aim: How do we do formula writing? Do Now: what is the difference between the empirical formula and the molecular formula? Homework: Text Book Pg 395 18-27

  43. 3. Structural Formulas • Structural formulas indicate the type, number, and arrangement of atoms within a compound. O O = C = O H H [Na]+ [Cl]-

  44. Glucose

  45. VIII. Formula Writing a)Ionic compounds: Given the name of the compound: 1. Determine if the compound is binary or ternary. 2. If the compound is binary, write out the symbols for the two elements present: metal then non-metal. * Elements with lower electronegativity values are usually listed first in a chemical formula. 3. Find and write out the oxidation number for the metal and non-metal. 4. Perform the crisscross method to determine the subscripts. * Eliminate the signs from the subscripts and if necessary reduce the subscripts to the lower whole number.

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