Atoms and Bonding

1. Atoms and Bonding Chapter 4 Unit – Introduction to the chemistry Physical ScienceMrs. N. Castro

2. Homework • Work chapter’s vocabulary words in your home. • Remember, as always, do it in the index cards. • Front part – the vocabulary word. • Back part – vocabulary word definition. • Bring all the index card with a rubber band for tomorrow. • Value – 10 points. All the vocabulary words.

3. Atoms, bonding and the Periodic Table • Read pages 125 to 127. • Answer the following: • What is compound? • How the elements are combined? • What is chemical bond? • Which particle from the atom can form chemical bond? • Compound – Substance made of two or more elements chemically combined in a specific ratio or proportion. • The elements are combined using a chemical bond. • Chemical bond – force of attraction that holds atoms together as a result of the rearrangement of the electrons between them. • Electrons

4. Atoms, bonding and the Periodic Table • Is there all electrons in the same energy levels? • Which electrons can form chemical bond, all of them or any specific? • In which energy level we can find the valence electrons? • What determines the valence electrons in each atoms? • No • Valence electrons • In the higher energy level with the higher amount of energy. • They determines the chemical property, to whom the atom will establish a chemical bond and how will be.

5. Atoms, bonding and the Periodic Table Which is the maximum of chemical bonds that the atom can do? Which family is able to perform the maximum chemical bond? Which is the minimum? Note: Families 3 to 12 will be study in future grades.

6. Atoms, bonding and the Periodic Table • How the valence electrons can help me to determines the quantity of chemical bonds the atom can perform? • Using the electron dot diagram or Lewis diagram. • To perform electron dot diagram: • Write the element symbol. • Surround the symbol by dots. • Each dot mean one valence electron. • Maximum numbers of dots – 8 • Maximum dots per side – 2 • Side - 4

7. Atoms, bonding and the Periodic Table • Atoms tend to be more stable if they have 8 valence electrons. • The only column in the Periodic Table with 8 valence electrons is family # 18. They are the most stable elements in the Periodic Table. This is the reason to explain why they don’t perform chemical bond with any element. • Family # 18 is nonreactive or stable. • He (Helium) is an exception, because it has only 2 valence electrons and it stable or non reactive element. It is part of the Family 18. • Another exception is H (Hydrogen), need only two valence electron to be stable. Actually, it has only one valence electron. Need only one more to complete it stability or nonreactive state.

8. Atoms, bonding and the Periodic Table Let’s work with Electron Dot Diagram. Teacher will perform some and you will perform more.

9. Atoms, bonding and the Periodic Table • Ways to forms atoms bonds: • Valence electrons may be transferred from one atom to another. One atom loose the valence electron and the other gain it. • Atoms share valence electrons. • Depend on how the atom form the bond, it will be the chemical bond. • Types of chemical bonds; • Ionic bonds – metal with nonmetal • Covalent bonds – nonmetal with nonmetal • Metallic bonds – metal with metal

10. Atoms, bonding and the Periodic Table • Metal • React by loosing valence electrons. • Reactivity will depend on how easily its atoms loose valence electrons. • Nonmetals • React when they gain or share enough electrons to have 8 valence electrons (except H). • Nonmetals usually combine with metals by gaining electrons. • Nonmetals can also combine with other nonmetals and metalloids by sharing electrons.

11. Atoms, bonding and the Periodic Table • Metalloids • Can either lose or share electrons when they combine with other elements. • Hydrogen • Its has 1 valence electron – Family 1 • Is nonmetal. • Properties very different from alkali metals. • Share its electron when form compound with other nonmetals to obtain a stable arrangement of 2 electrons.

12. Atoms, bonding and the Periodic Table Practice N • Element name? • Find the protons, electrons and neutrons. • Is N reactive or stable? • Mention two elements with properties similar to N. Explain why they are similar. • How many bonds N can performed? • With which elements, N will prefer to establish chemical bond? Why?

13. Atoms, bonding and the Periodic Table More practice Perform all exercises from your book, pages 125 to 129. Include “Assess Your Understanding”

14. Ionic Bonds Lesson 2 Pages 131 to 137

15. Ionic Bonds • To perform ionic bonds you need ions • ION is an atom or group of atoms that has an electric charge. Example: Cl-, Na+ • Types of ions: • Negative ion – when atom gain an electron. • Positive ion – when atom lose an electron. Examples: K = 1 valence electron, Family 1, lose the electron = K+ Br = 7 valence electrons, Family 17, gain 1 = Br _ Metals atoms are likely to lose electrons. Nonmetals atoms likely to gain electrons. Metal Nonmetal

16. Ionic Bonds • Common Ions • Monoatomic – ion made of one element • Example: K+ • Polyatomic – ions made of more than one elements. Can be positive or negative charges. • Example: HCO3- = one Hydrogen, one Carbon and three Oxygen = one negative charge • See table on page 132.

17. Ionic Bonds • Ionic bond – is the attraction between two oppositely charged ions. • Ionic compound – resulting compound from ionic bond. The ionic compound always is neutral, charge = 0. The total positive charge of all positive ions equals to total negative charge of all the negative ions. Let see some examples of ionic bonds.

18. Ionic Bonds Video

19. Ionic Bonds – Formulas • What is the different between this? K , KBr , K+ , Br- K = element symbol KBr = chemical formula K+ = positive ion Br- = negative ion • Chemical formula - is a group of symbols that show the ratio of elements in a compound. KBr – ratio = one K and one Br MgCl2 – ratio = one Mg and two Cl • The formula tell you: • The elements that component the compound. • The ratio of each element in the compound (subscripts).

20. Ionic Bonds – Formulas • Rules: • To write the formula for compound, always write the positive charge first and then the negative charge. • Add subscripts that are needed to balance the charge. ONLY SUBSCRIPTS. • If no subscripts is written, its is understood that the subscript is 1. Example: NaCl = no subscripts for both elements; means: one Na and one Cl. • Formulas for compounds off polyatomic ions are written in a similar way.

21. Ionic Bonds – Names • Rules for ionic compounds: • The name of the positive charge comes first, followed by the name of the negative ion. NaCl – Sodium chloride MgBr – Magnesium bromide • The name of the positive ion is usually the name of a metal. • If the positive charge is a polyatomic ion, use the name of the polyatomic ion. NH4Br – Ammonium chloride • If the negative ion is a single element, the end of its name changes to –ide. MgO – Magnesium oxide • If the negative ion is polyatomic, its name usually ends in –ate or –ite. NH4NO3 – Ammonium nitrate. NH4HCO3 –Ammonium bicarbonate

22. Ionic Bonds • Exercises: • Teacher exercises • Find the ratio of each element in the compound. • Write chemical formula for ionic's compounds. • Name ionic's compounds. • Complete exercises from pages 131 to 135.

23. Ionic Bonds • Properties of ionic compounds: • Hard. • Brittle crystals. • Forms solids by building up repeating patterns of ions. This three dimensional arrangement is called crystals. • High melting points • When dissolved in water or melted, they conduct electric current. • Attraction between ions is very strong, so it takes a lot of energy to break the bond. • When ionic compounds dissolve in water, they become ions again. • The charged particles carry electric current where the neutral ionic crystals do not.

24. Ionic Bonds • Extra exercises: • A fluorine (F) ion has a charge of 1-. An aluminum (Al) ion has a charge of 3+. • Draw for both atom the electron dot diagram. • Explain how the F and Al would exchange valence electrons to form an ionic compound. Draw the explanation first and then explain in your own words. • Write the compound’s chemical formula. • Name the compound. • A potassium ion has a charge of 1+. A sulfide ion has a charge of 2-. What is the chemical formula for ionic compound and the name. • Name the following compound MgO • Perform exercises on pages 136 and 137.

25. Covalent Bonds Lesson 3 Pages 138 to 145

26. Covalent Bonds • It formed when 2 atoms share electrons. • Usually are between nonmetals atoms. • The attractions between the shared electrons and the protons in the nucleus of each atom hold the atom together in a covalent bond. • The product of a covalent bond is called a molecule. • A molecule is a neutral group of atoms joined by covalent bonds.

27. Covalent Bonds

28. Covalent Bonds • Exercise: • Draw the Lewis diagram for Bromine atom. • Draw the Lewis diagram for the Br2 molecule. • Perform exercise on page 139.

29. Covalent Bonds • Atoms can form single, double and triple covalent bonds by sharing one or more pairs of electrons. • Double bond – 2 atoms share 2 pairs of electrons

30. Covalent Bonds • Triple bond – share 3 pairs of electrons.

31. Covalent Bonds • Identify the single, double and triple covalent bonds: Draw the Lewis diagram for (CH3)2SO - dimethylsulfoxide

32. Covalent Bonds • Solution C H S O Fam. 14 Fam.1 Fam.16 Fam.16

33. Covalent Bonds • Properties: • Do not conduct electric current when melted or dissolved in water. • Lower melting point and boiling point.

34. Covalent Bonds • Sometimes, atoms of some elements pull more strongly on the shared electrons of a covalent bond than do atoms of other elements. As a result, the electrons ser shared unequally. • Unequal sharing of electrons causes covalently bonded atoms to have slight electric charges.

35. Covalent Bonds • Types of covalent bonds: • Non polar bond – a covalent bond in which electrons are shared equally. Example: H2

36. Covalent Bonds • Polar bonds – Covalent bonds in which electrons are share unequally. When electrons in a covalent bond are shared unequally, the atom with the stronger pull gains a slightly negative charge. The atom with the weaker pull gains a slightly positive charge. Example: HF

37. Covalent Bonds video

38. Bonding in Metals Lesson 4 Pages 146 -151

39. Bonding in metals • Metal atoms lose electrons easily because they do not hold their valence electrons very strongly. • The loosely held valence electrons in metal atoms result in a type of bonding that happens in metals. • Most metals are crystalline solids. • A metal crystal is composed of closely packed, positively charged ions. • Each metal ion is held in the crystal by a metallic bonds. • Metallic bonds is an attraction between a positive metal ion and electrons surrounding it.

40. Bonding in metals • Metallic bonds is formed an attraction within metal atoms. • Video ...\..\Clark-9805.Fig.4 (2).mov • Properties: • Luster – When the light strikes these valence electrons, they absorb the light and then re-emit the light. • Malleability and Ductility – The positive metal ions are attracted to the loose electrons all around them rather than to other metal ions. These ions can be made to change position. However, the metallic bonds between the ion and the surrounding electrons keep the metal ions from breaking apart from one another.

41. Bonding in metals • Thermal conductivity – Metals conduct heat easily because the valence electrons within a metal are free to move. • Electrical conductivity – Metals conduct electric current easily because the valence electrons in a metal can move freely among the atoms. • Alloys – is a mixture made of two or more elements, at least one of which is a metal. • Generally are stronger and less reactive than the pure metals from which they are made. • Example: stainless steel = iron +carbon + nickel + chromium • Gold jewelry = gold+ cooper + silver

42. Assessyourunderstanding Allbondings

43. All Bondings • http://www.bsc2.ehb-schweiz2.ch/Chemie/Simulationen%20Chemie/Bindung/Bindung%20Hundeanalogie.htm • Compare and contrast the following concepts:

44. END Complete chapter review