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Chemical Bonds

Chemical Bonds. Forming Chemical Bonds. The force that holds two atoms together is called a chemical bond . The valence electrons are the electrons involved in forming chemical bonds. Elements tend to react to acquire eight electrons. This is called a stable octet .

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Chemical Bonds

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  1. Chemical Bonds

  2. Forming Chemical Bonds • The force that holds two atoms together is calledachemical bond . • The valence electrons are the electrons involved in forming chemical bonds. • Elements tend to react to acquire eight electrons. This is called a stable octet. • Noble gases(group VIIIA/18) have this structure (octet) and are inert (does not form bonds). • Atoms can gain, lose, or share electronsto reach an octet.

  3. Forming Ions • Positive ions (cations) are formed when atoms lose one or more valence electrons. [Usually these atoms are metals.] • Reactivity of metals (cations) are based on how easily they lose electrons. [ Ionization energy] • Negative ions (anions) are formed when atoms gain one or more valence electrons. [These atoms tend to be non-metals.] • Electronegativity is the ability to attract or gain electrons.

  4. Forming Ionic Bonds • Ionic bonds: Complete transfer of valence electrons between atoms (difference of electronegativity of 1.6 or greater). • Bond between a metal and a nonmetal • Two neutral atoms will form ions. The resulting compound is called an ionic compound.

  5. Properties of Ionic Compounds • Form crystal lattice bonding structure. • Example: NaCl (sodium chloride) • High melting point and boiling point due to strong electrostatic charge between the atoms (cation is attracted to anion after the transfer of electrons). • Hard, rigid and brittle solids. • Conducts electricity in liquid state or dissolved in water only.

  6. Forming Covalent Bonds • Covalent bonds: Sharing of valence electrons between atoms (difference of electronegativity of less than 1.6). • Occurs usually between elements close to each other on the periodic table (mostly 2 non-metals). • The resulting compound is called a molecule.

  7. Properties of Covalent Compounds • Have definite and predictable shapes. • Low melting and boiling points. • Relatively soft solids. • Can exist as solids, liquids or gases.

  8. Metallic bonds • This weak bond is formed due to the attraction between kernels and the mobile electrons in a metal lattice. • Properties: • High thermal and electrical conductivity. • Luster and high reflectivity. (shiny) • Malleability and ductility. (They can be beaten or shaped without fracture.) • Variability of mechanical strengths

  9. Complete your Venn diagrams • Requirements: • 5 for ionic only • 5 for covalent only • 1 similarity

  10. Comparison of Bonding Types Ionic Covalent Metal and nonmetal ions 2 nonmetals molecules molten salts conductive non-conductive valence e- transfer of electrons sharing of electrons high mp & bp low bp & mp DEN < 1.6 DEN > 1.6

  11. Ionic and Covalent Song https://www.youtube.com/watch?v=5Zge2KtT8l8

  12. Diatomic Elements • These atoms are never found alone in nature. They are always bonded to something. (that can mean another atom of their same element) • Memorize: H2, N2, O2, F2, Cl2, Br2, I2 • To help memorize remember at 7 make a 7 + H.

  13. The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together. Subtract to find the difference.

  14. Examples in calculating the difference of electronegativity • HF, we would subtract the electronegativity of hydrogen (2.1) fromfluorine (4.0). • 4.0 - 2.1 = 1.9 • O2 is a diatomic element. Since the two oxygen's have the same electronegativity, the difference between them is 0. • MgCl2 is an ionic compound. We would subtract the electronegativity of magnesium (1.2) from chlorine (3.0). • 3.0 - 1.2 = 1.8

  15. Electronegativity and Bond Character • The character and type of a chemical bond can be predicted using the electronegativity difference of the elements that are bonded. • Electronegativity is a measurement of how strong the atom is pulling on electrons that it is sharing in a bond with another atom. • For identical atoms, the electronegativity difference is zero, the electrons are shared equallyand the bond is considered nonpolar covalent, which is a pure covalent bond.

  16. A system of “drawing” bonds Shows how valence electrons are arranged Dots represent valence electrons Pairs of dots represent bonding pairs of electrons. Show the reaction between sodium and chlorine: Show the reaction between calcium and bromine: Lewis Dot Structures Ionic

  17. Determining Number of Covalent Bonds • To determine how many bonds exist in a molecule, use the following formula N - A = # bonds 2 -- Where N is the # of needed electrons, which is 8 for all elements but H, which is 2. -- Where A is the # of available electrons, which is the number of valence electrons.

  18. Drawing Lewis Structures Covalent 1.Determine the # of bonds needed 2.Draw a skeleton structure. 3.Connecttheatoms with the # of bonds. (determined in step 1) 4.Finish by making sure all atoms in the structure have an octetof electrons. 5.Remember, if the structure has a positivecharge, it has fewer available electrons, and if it has a negative charge it has more available electrons.

  19. Lewis Structure Practice • Draw Lewis Dot Structures for the following molecules and polyatomics: CO2 N2 O2 Cl2 H2O SO42-

  20. Exceptions to the Octet Rule • There are two elements which don’t want an octet of electrons and those are Be (wants 4) and B (wants 6). • Third-rowand heavier elements often satisfy the octet rule but can exceed the octet rule using their empty valence d orbitals, especially when surrounded by highly electronegative atoms such as chlorine, fluorine, bromine, and oxygen.

  21. Exceptions to the Octet Rule • The exceptions are easy to recognize-- when you calculate the # of bonds, the formula will give you an answer that makes no sense.If this happens, do the following: • Connect all of the atoms to the central element with one bond. • Give all of the surrounding elements an octet of electrons. • Count the number of electrons you have put in the structure so far. - If it equals the available # of electrons you calculated, you’re done. - If it is less than the available # of electrons you calculates, place the needed number of electrons around the central element.

  22. Exceptions to the Octet Rule • Draw Lewis Dot Structures for the following molecules: BF3 PCl5 SF6 XeF2

  23. Carbon Bonding • Carbon follows the rules previously discussed, but it has a few unique properties. • Carbon is always the central atom in its compounds and it tends to bond to itself. • Carbon always has four bonds. • Anytime you have a choice for a skeleton structure, always go for the most symmetrical option.

  24. Carbon Bonding • Draw Lewis Dot Structures for the following molecules: C2H6 C3H8 C2H2

  25. Bond Strength and Length • C C C C C C Bond Type single double triple Bond Length 154 134 121 Bond Strength 414 611 837 Bond lengths are in picometers and bond energies are in kJ • Stronger bonds = higher boiling & melting points • Weaker bonds = lower boiling & melting points

  26. Resonance • Resonance occurs when you have a double bond which can be placed in more than one location. • More than one valid Lewis structure can be written for a molecule. Examples: O3, NO3-, NO2-, SO22-, CO32-

  27. Let’s Practice! • H2 • SO2 • CCl4 • SiH4 • PCl5 • XeCl6 • H2O • SO4-2 • NO3- • C2H3O2-

  28. Molecular Shape • Valence Share Electron-Pair Repulsion(VSEPR) model allows us to predict the molecular shape by assuming that the repulsive forces of electron pairs cause them to be as far apart as possible from each other. • Only the valence electron pairs are considered in determining the geometry.

  29. 2 charge clouds, linear 3 charge clouds, trigonal planar 4 charge clouds, tetrahedral Effect of the number of electron pairs around the central atom

  30. PREDICTING EXPECTED GEOMETRY ACCORDING TO VSEPR THEORY • Lewis dot structure determines the total # of electrons around the central atom. • Multiple bonds (double and triple) count as one. • The number of bonding and nonbondingelectron pairs around the central atom determines the geometry of electron pairs and the molecular geometry. • Lone e pairs affect geometry more than bonding pairs. • The shape is referred to asbentif there are lone pairs on the central atom.

  31. Molecular Shapes 2,3,4 Electron Pairs

  32. Molecular Shape Practice • Example: • H2CO • HCN

  33. Molecular Shape Practice • Example: • CO2 • PF3

  34. Molecular Shape Practice • Example: • SO42-

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