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Chemical Bonding

Chemical Bonding. Quarter 2 Unit 4. Compounds and Atomic Stability: Learning Objectives. Understand why atoms form compounds Understand chemical stability and the octet rule Understand how Ionic bonds and Ionic compounds are formed. How Elements Form Compounds.

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Chemical Bonding

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  1. Chemical Bonding Quarter 2 Unit 4

  2. Compounds and Atomic Stability: Learning Objectives • Understand why atoms form compounds • Understand chemical stability and the octet rule • Understand how Ionic bonds and Ionic compounds are formed

  3. How Elements Form Compounds • Compounds:a chemical combination of two or more different elements joined together in a fixed proportion • Collisions between particles of the atom cause reactions • Reactions between atoms Involve Only The Electron Cloud! • Remember, chemical properties of elements on the periodic table repeat because the pattern of valence electrons repeat in each period • Valence electrons of colliding atoms react to form compounds • Compounds form when electrons in atoms rearrange to achieve stable electron configurations.

  4. Chemical Stability • Nobel Gasses: they are almost completely un-reactive • None of these have ever been found naturally in the environment as a compound. • They are extremely un-reactive or stable • Group 18 (except He) all have 8 valence electrons • Electron arrangement determines chemical properties • The electron arrangement of the noble gasses is the cause of their stability (extremely un-reactive)

  5. Check Your Understanding Turn and discuss the following with the person on your right, then answer each of the questions in your notebook. • Identify each of the following as a compound or not a compound • Water • Nitrogen • Carbon dioxide • Deuterium • What subatomic particles are involved in forming compounds? • Why do elements form compounds? • Which group of elements rarely forms compounds? Explain why.

  6. Answers • Identify each of the following as a compound or not a compound • Water--- compound • Nitrogen---- no, an element • Carbon dioxide-----compound • Deuterium-----no , an element • What subatomic particles are involved in forming compounds?--- valence electrons • Why do elements form compounds? --To rearrange electrons in order to achieve a stable electron configuration • Which group of elements rarely forms compounds? Explain why. ---Noble gasses, because they already have a stable electron configuration

  7. The Octet Rule • Atoms combine to become more stable • Atoms become stable by having eight electrons in their outer energy levels (He is an exception w/ 2 because it is so small) • Atoms become stable by achieving a noble gas configuration

  8. Achieving Chemical Stability • Collisions between atoms, which involve enough energy, can cause valence electron rearrangements • Forming a stable octet • Noble gas configuration • Total number of electrons never change

  9. Establishing Stable Octets • 2 options • Transfer of electrons between atoms • Sharing electrons between atoms

  10. Electrons are Transferred Example • Na- sodium is a shiny gray metal which quickly oxidizes upon exposure to the atmosphere • Cl - Chlorine gas is a dense, pale yellowish-green, poisonous, gas • Na and Cl react to form–salt

  11. Electron Transfer On a Subatomic Level • Na- a group 1 element, with 1 valence electron • Cl - a group 17 element, with 7 valence electrons • How can the valence electrons be rearranged to provide a stable octet for each?

  12. Chlorine Gets a Stable Octet • If the Cl atom gains an electron from the Na atom, it will achieve a noble gas configuration • It will be stable • It will have a complete octet • It will also have a negative charge, because it has an extra electron • It will be an ion

  13. Na Gets a Stable Octet • If Na gives away its 1 valence electron, what will it have for a valence number? • 11 electrons minus 1 electron = 10 electrons • 1s22s22p6 - 8 valence electrons • It will be stable • It will have a complete octet • It will also have a positive charge, because it has one less electron • It will be an ion

  14. Transfer of the Electron Forms an Ionic Bond

  15. An Ionic Bond • The exchange of the electron creates 2 ions, • 1 positively and 1 negatively charged • Ionic bond: A strong attractive force between atoms with opposite charges which is formed by an electron transfer between atoms. • Ionic bonds form ionic compounds • A compound that is made up of ions • Na+ Ions are attracted to all nearby Cl- Ions and visa versa thereby forming crystal structures • Cl- • Na+

  16. The Results of Ionic Attraction • Affects properties of the compound • Example; sodium chloride (salt) is a crystal because of intermolecular forces of attraction between ions, it is a solid at room temperature • Melting Ionic Compounds: Breaking the strong crystal structure requires a lot of energy, therefore the melting point of NaCl is more than 800°C • Hardness and Brittleness: It takes a great deal of force to break the structure of an ionic crystal

  17. Naming Binary Ionic Compounds • Binary compounds are compounds with only two different elements. • First write the name of the positively charged ion (usually a metal) • Then add the name of the negatively charged ion (nonmetal) • Modify the negatively charged ion name to end in -ide • Example NaCl is called sodium chloride NOT sodium chlorine

  18. Check Your Understanding • Name compounds formed by the following ions: • Mg+ Cl- • Cl- Ca+ • Cl- K+ Name these compounds: • AlN • KI • ZnO

  19. Answers • Name compounds formed by the following ions: • Mg+ Cl- Magnesium chloride • Cl- Ca+ calcium chloride • Cl- K+ potassium chloride Name these compounds • AlN aluminum nitride • KI potassium iodide • ZnO zinc oxide

  20. Oxidation Numbers • Charged atoms or compounds are called Ions • The total charge on the ion is known as the Oxidation Number of the atom • Examples: • Mg+2 is magnesium ion, the charge is positive 2, the oxidation number is 2+ • F- has an oxidation number of 1- • Some metals have the same oxidation number in all compounds (memorize this) • Group 1 elements , oxidation number = 1+ • Group 2 elements , oxidation number = 2+ • Aluminum, oxidation number = 3+

  21. The Charge of Ionic Compounds • In Ionic Compounds, the total positive charge is equal to the total negative charge • One Mg2+ ion will combine with 2 Cl- ions • Forming MgCl2, The total positive charge is 2+, the total negative charge is 2- • +2+-2 = 0 • In a correctly written formula, the sum of the total positive charges and the total negative charges = 0

  22. Predicting Oxidation Numbers • Oxidation numbers for most elements can be predicted from their position on the periodic table. • Groups 3-12 Transition metals are difficult as many of these elements Have more than one oxidation number depending on the reaction • Group 13 elements have 3 valence electrons, so will lose 3 and have an oxidation number of 3+ • Group 14 may have 2+ or 4+ oxidation number • Groups 15, 16 and 17 tend to gain electrons to complete the octet since they are already ½ full • Their oxidation numbers are 3-, 2-, and 1- respectively • They can also lose electrons and have positive oxidation numbers • The tendency to lose electrons increases as you move down the column

  23. Check for Understanding • Predict the oxidation numbers for the following elements: • Al • N • Cl • Mg • S • Na • K • O • Ga • P • Se • Br

  24. Answers • Predict the oxidation numbers for the following elements: • Al 3+ • N 3- • Cl 1- • Mg 2+ • S 2- • Na 1+ • K 1+ • O 2- • Ga 3+ • P 3- • Se 2- • Br 1-

  25. Writing Chemical Formulas For Ionic Compounds • The key to writing formulas is to make the oxidation numbers add to zero, making a neutral compound. • Example: • Ca 2+ , located in group 2 • F 1-, located in group 17 • The formula for a compound of these elements is • CaF2 1(2+) + 2( 1-) =0 The compound is neutral

  26. Representing Compounds as Formulas • The formula of a compound tells: • what elements make up the compound • and how many of each element are present in one unit of the compound • Example : H2O two H for each O

  27. Progress Check • Write formulas for the following compounds: • Sodium Fluoride • Potassium Chloride • Rubidium bromide • Sodium selenide • Potassium oxide • Lithium sulfide • Strontium fluoride • Calcium Chloride

  28. Answers • Write formulas for the following compounds: • Sodium Fluoride (Na+1 F-1) NaF • Potassium Chloride (K+1 Cl -1) KCl • Rubidium bromide (Rb+1 Br-1) RbBr • Sodium selenide (Na +1 Se -2) Na2Se • Potassium oxide (K+1 O-2) K2O • Lithium sulfide (Li+1 S -2) Li2S • Strontium fluoride (Sr 2+ F 1-) SrF2 • Calcium Chloride (Ca 2+ Cl 1-) CaCl2

  29. The Formation Of Ionic Compounds • Lab Learning Objective: To model the transfer of electrons thereby achieving noble gas configurations and the formation of stable ionic compounds • Apply what you have learned! • The Formation Of Ionic Compounds (MiniLab 4.2) • Question: What other atoms give up and gain electrons (creating ions, forming ionic bonds) to form ionic compounds?

  30. Pre-Lab Procedure • You and your lab partner must locate the following atoms on the periodic table: Li, S, Mg, O, Ca, N, Al and I. • Using the information on the Periodic Table, and what you have learned, complete the table with the Valence electrons, class of element, Lewis dot diagram, and electron configuration for the given atoms • Based on what you have discovered about the atoms above, construct a hypothesis predicting which of the atoms will give up electron(s) and which will receive electron(s) when forming compounds. • REMEMBER a hypothesis is an if… then… statement which answers the lab question • Have your hypothesis checked before you begin the lab.

  31. The Formation Of Ionic Compounds • Learning Objectives: • To model the transfer of electrons thereby achieving noble gas configurations and the formation of stable ionic compounds • To name ionic compounds • To identify formulas of ionic compounds • Identify a pattern which can be used to predict which atoms will form Ionic compounds • Follow the Lab instructions to complete your investigation. • Prepare for a poster presentation

  32. Polyatomic Ions • Some Ions are formed by more than one or two types of atoms • They are called polyatomic ions • Some of these are very common, so recognizing them will be very helpful to you. • Let’s look at some • Some do not follow simple naming rules, often because they were named before rules for naming were written • These we need to use a reference sheet for, or memorize! • Use your reference sheets to help name complex polyatomic ionic compounds

  33. Conventions for Naming Polyatomic Ionic Compounds When Given the Formula • Steps • Name positive ion (cation) first, then negative ion (anion) second. • If one of these ions has more than 1 atom in it, look up the name on the chart • You should plan on memorizing the ones I have starred! • Determine if you need a Roman numeral in the name • If cation is NOT a transition metal, then NO Roman numeral • If the cation is a transition metal then see if it can have more than 1 oxidation number, if not then go to step 4 • If yes then go to step 3 • Determining the Roman numeral, usually the number of anions = the charge on the cation and the number of cations =anions. • Example: Fe2[SO4}3 • There are three SO4 sulfate ions -- charge of 2-= 6 - • There are two Fe atoms – charge of ? = must total 6 • Fe has to be 3+ so use Roman Numeral III Fe(III) • Name is Iron (III) sulfate! • Check to see that the sum of oxidation numbers= 0 • If yes then correct! • Example : • 3, SO4 ions = 3 x 2- = 6- • 2 Fe(III) ions = 2 x 3+ = 6+ • 6- + 6+ = 0 !!!!!

  34. Let’s Try Another One • NH4Cl • Look at your list of common polyatomic ions, do you see any here? • Yes NH4 is ammonium it is a cation – 1+ • Follow step 1--- • Ammonium is the first part of the name • Chlorine is the second atom, the anion- 1-, becomes chloride (just like before) • So far we have ammonium chloride • Follow step 2--- • Is ammonium a transition metal? No! • Follow step 4 • 1+ + 1- = 0 correct!  ammonium chloride

  35. Now you try it alone! • Name the following : • Fe(NO3)3

  36. Solution • Iron(III)nitrate • Iron  cation = 2+ or 3+ • NO3  nitrate = 1- • Three anions, so the charge needs to be 3+ so Roman numeral is III • Check • 1 Fe(III) has charge of 3+ • 3 NO3 ions have a charge of 3 x 1- = 3- • 3+ + 3- = 0 correct!  Iron(III)nitrate

  37. Let’s Try Another • Al(CN)3

  38. Answer • Al(CN)3 • Aluminum Cyanide

  39. Establishing Stable Octets • 2 Options • Transfer of electrons between atoms – Form Ionic Bonds • Sharing electrons between atoms

  40. Obtaining a Stable Octet by Sharing • Example: • H (hydrogen) • A gas at room temperature • Has one valence Electron • O (oxygen) • A gas at room temperature • A group 16 element • Has 6 valence electrons • Could they achieve a stable octet by transferring electrons? • Could H give up its only electron and have a noble gas configuration? • No electrons!

  41. Sharing Electrons Between Atoms • When atoms collide with enough energy to cause a reaction • And neither atom attracts electrons strongly enough to take electrons from the other atom (small difference in electronegativity) • The atoms combine by sharing valence electrons

  42. Reaction of Hydrogen and Oxygen • Hydrogen needs one more electron to have the same electron configuration as He • Oxygen needs two more electrons to have the same electron configuration as Ne • Hydrogen and Oxygen can share one electron from each atom. • This makes Hydrogen stable • But Oxygen still has only 7 valence electrons

  43. Sharing to Make Oxygen Stable • Oxygen gets a complete octet by sharing an electron with another hydrogen • This explains the formula of H2O

  44. Electrons always rearrange in a chemical reaction

  45. Covalent Bond • The attraction of two atoms for a shared pair of electrons is called a covalent bond • Notice, neither atom has an ionic charge • A compound whose atoms are held together by covalent bonds is called a covalent compound • A molecule is an uncharged group of two or more atoms held together by covalent bonds

  46. Sharing More Than Two Electrons • More than two electrons can be shared. • The reaction Between Carbon and Oxygen for example • You can arrange these 16 valence electrons to produce a molecule in which all three atoms have a noble gas configuration (a complete octet)

  47. Sharing two pairs of Electrons: Double Covalent bonds are formed Atoms can also form Triple bonds

  48. Molecular Elements • Molecular Elements: a molecule that forms when atoms of the same element bond together • They are not compounds (not two elements) 7 nonmetal elements naturally found as diatomic molecules: • Hydrogen • Nitrogen • Oxygen Gases • Fluorine • Chlorine • Bromine -liquid • Iodine- solid Allotropes: molecules of an element that form different crystalline structures. • The properties of allotropes are usually different, because structure is important

  49. Formulas and Names of Covalent Compounds • Binary inorganic compounds (2 elements, not carbon) • The Suffix –ide • Write first nonmetal followed by the name of the second nonmetal with its ending changed to – ide • The element closest to the left of the periodic table is written first ( some exceptions with H) • If both are in the same group, name the one lower on the column 1st. • Indicating the number of atoms • Add a prefix to the name of each element indicating the number of each element

  50. Add prefix to indicate number of atoms • If only one atom of the first element is listed, the mono is usually left out • If adding the prefix creates double vowels, the first is usually omitted

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