CHEMISTRY 111 WEEK-5-6-PERIODIC-TABLE.pptx

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1. WEEK 5 - 6 Periodic Table of Elements Classifications of the Elements Trends in the Periodic Table of Elements Compounds and Chemical Formula

2. At the end of the week 6, the student will be able to: Identify Group and Period where the element belongs. Predict the atomic, physical and chemical properties of elements using Trends in the periodic table. Name the given chemical formula for simple ionic and molecular compounds.

3. I. Periodic Table of Elements: Classification of Elements Quick Trivia: 118 Elements 94 elements: naturally-occurring 24 elements: man-made

4. Three Main Categories of Elements

5. Three Main Categories of Elements METALS Majority of elements in the PT Lies on LEFT side of periodic table.

6. Three Main Categories of Elements Metals include: Alkali metal Alkaline earth Transition metal Basic metal Lanthanide, and Actinide groups. These elements have metallic character , which means atoms easily lose electrons .

7. Three Main Categories of Elements Metals display the following properties: Solid at room temperature M etallic luster High melting point Good conductor of heat Good conductor of electricity Malleable - able to be pounded into sheets Ductile - can be pulled into wire High density (exceptions: lithium, potassium and sodium) Corrode in air or seawater Lose electrons in reactions

8. Three Main Categories of Elements II. NONMETALS Concentrated on the RIGHT side of PT

9. Three Main Categories of Elements Nonmetals include: Nonmetals element group Halogens Noble gases

10. Three Main Categories of Elements Nonmetals display the following properties: Dull – not shiny Poor conductor of heat Poor conductor of electricity B rittle – break or shatter easily Lower density Lower melting point and boiling points Gain electrons in reactions

11. Three Main Categories of Elements III. METALLOIDS Semimetals Zigzag/STAIRCASE pattern on periodic table Same physical appearance as metals but behaves chemically like non-metals.

12. Three Main Categories of Elements III. METALLOIDS Semimetals B Si Ge As Sb Te Po

13. Three Main Categories of Elements Nonmetals display the following properties: Could be dull or shiny Conduct heat and electricity, but not as well as metals Good semiconductors Usually malleable Usually ductile Can both gain and lose electrons in reactions

14. CLASSIFICATION OF ELEMETS

15. CLASSIFICATION OF ELEMETS 1. Representative Elements 2. Transition Elements 3. Inner Transition Elements

16. CLASSIFICATION OF ELEMETS 1 . Representative Elements Groups 1A – 8A Four Groups with SPECIAL NAMES : Alkali Metals – Group 1A Alkaline-Earth Metals – Group 2A Halogens – Group 7A Noble Gas – Group 8A

17. CLASSIFICATION OF ELEMETS 2. Transition Elements The group designated by the numeral number and the letter B. Ex. Iron, Copper, Gold

18. CLASSIFICATION OF ELEMETS

19. CLASSIFICATION OF ELEMETS 3.Inner Transition Elements Lanthanides (Lan-Ce- Lut ) Cerium (Ce) – Lutetium (Lu) 58 – 71 b. Actinides (Ac-Tho-La) Thorium (Th) – Lawrencium (Lr) 90 - 103

20. GROUP OF ELEMETS GROUP 1: ALKALI METALS Total: 7 Elements: H, Li, Na, K, Rb, Cs, Fr Characteristics: Except HYDROGEN 1 Valence Electron Active Metals – readily react Abundant: Sodium and Potassium Rare: Li , Rb , Cs All are SOFT Reactive with Oxygen and moisture

21. GROUP OF ELEMETS Alkali Metal Compounds: Table Salt – ( NaCl or Sodium chloride ) Potash – ( K 2 CO 3 – Potassium carbonate ) Washing Soda – ( Na 2 CO 3 – Sodium carbonate ) Lye – ( NaOH – Sodium Hydroxide ) Baking Soda – ( NaHCO 3 ) – Sodium Hydrogen Carbonate or Sodium Bicarbonate)

22. GROUP OF ELEMETS GROUP 2A: ALKALINE-Earth METALS Total: 6 Elements: Be, Mg, Ca, Sr, Ba, Ra Characteristics: 2 Valence Electron Active metals but not as reactive as Alkali Higher Melting Point Harder and Stronger than Alkali

23. GROUP OF ELEMETS Alkaline-Earth Metal Compounds: Milk of Magnesia – Mg(OH) 2 – Magnesium hydroxide Limestone – CaCO3 – Calcium Carbonate Calcium Phosphate – Ca3(PO4)2 : used by vertebrates in formation of bones and teeth

24. GROUP OF ELEMETS GROUP 3A: BORON GROUP Total: 6 Elements: B, Al, Ga, In, Tl, Nh Characteristics: 3 Valence Electron 1 metalloid, 5 metals Reactive Solids at room temperature

25. GROUP OF ELEMETS GROUP 4A: CARBON GROUP Total: 6 Elements: C, Si, Ge, Sn, Pb, Fl Characteristics: 4 Valence Electron 1 nonmetal, 2 metalloid, 3 metals Each element in this group has atoms that can GAIN, LOSE, or SHARE 4 electrons when reacting with other elements.

26. GROUP OF ELEMETS GROUP 5A: NITROGEN GROUP Total: 6 Elements: N, P, As, Sb, Bi, Mc Characteristics: 5 Valence Electron 2 nonmetal, Nitrogen and Phosphorus 2 metalloids, As and Sb 1 metal, Bi

27. GROUP OF ELEMETS GROUP 6A: OXYGEN GROUP Total: 6 Elements: O, S, Se, Te , Po, Lv Characteristics: 6 Valence Electron 3 nonmetals: O, S, Se Reactive Solid at room temperature (Except Oxygen)

28. GROUP OF ELEMETS GROUP 7A: HALOGEN GROUP Total: 6 Elements: F, Cl, Br, I, At, Ts Characteristics: 7 Valence Electron Most reactive nonmetals Have strong tendency to GAIN “ one ” more electron Active nonmetals Present in nature only I the form of their compounds ALL FORM COMPOUNDS with Na (Sodium)

29. GROUP OF ELEMETS GROUP 8A: NOBLE GASES Total: 7 Elements: He, Ne, Ar , Kr, Xe, Rn, Og Characteristics: 8 Valence Electron ALL exist as Monoatomic species Nonmetals Inert gases Stable

30. TRENDS IN PERIODIC TABLE

31. PERIOD VS GROUP

32. PERIOD VS GROUP

33. TRENDS IN PERIODIC TABLE M- metallic Character E- electron Configuration and Valence Electron A- atomic Size I- ionization Energy E- electronegativity

34. 1. Metallic Character Metals - elements whose atoms tend to LOSE electrons during chemical reactions. Nonmetals - elements whose atoms tend to GAIN electrons during chemical reactions.

35. TREND METALLIC CHARACTER Decreases from LEFT to RIGHT (across period) Increases from TOP to BOTTOM (down the group) NONMETALLIC CHARACTER Increases from LEFT to RIGHT (across period) Decreases from TOP to BOTTOM (down the group)

36. Most metallic natural element is Cesium . While the most metallic element of all is Francium . The least metallic or most non-metallic element is Fluorine .

37. Metallic character increases moving down a periodic table group and decreases moving across a period

38. Elements in the same Group have the same number of valence electrons 2. Electron Configuration and Valence Electron Chemical reactivity of the atoms depends on Electronic Configuration or the order of the electrons in the energy levels in their atoms. Valence shell - Outermost shell of atom Valence Electrons - Involved in the forming of chemical bonds

39. 2. Electron Configuration and Valence Electron

40. 2. Electron Configuration and Valence Electron Number of electrons in an atom of a given elements is the same as the element’s atomic number . Period Number is the same as the Number of Shells . For Representative Elements, the number of Valence electrons is the SAME as the Group Number. For Representative Elements, the number of Valence Electrons INCREASES by one as you proceed across a given period.

41. 3. Atomic Size Atomic size INCREASES down a Group. Atomic Size DECREASES across the Period.

42. 3. Atomic Size As the charge on the nucleus increases (more protons) without adding an additional shell of electrons, the outer electrons are tight bound, thus decreasing the atomic radius from LEFT to Right across the Period.

43. 4. Ionization Energy - the amount of energy that it takes to remove an electron from an atom. INCREASES – across Period DECREASES – down the Group Group 1A – lowest ionization energies. - Giver of electrons

44. 5. Electronegativity - the atom’s attraction for the electrons it shares in chemical bond with another atom. INCREASES – across Period DECREASES – down the Group

45. Electron Affinity - the change in energy of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion (Anion). INCREASES – across Period DECREASES – down the Group

46. Compounds and Chemical Formula READY INYONG BRAIN.

47. II. Introduction to Compounds Substance made from two or more different elements that are chemically bonded together in fixed proportions. Can be separated by chemical means. Examples: Water (H₂O): A compound of hydrogen and oxygen. Carbon Dioxide (CO₂): A compound of carbon and oxygen. Sodium Chloride (NaCl): A compound of sodium and chlorine.

48. II. Types of Compounds IONIC COMPOUNDS COVALENT COMPOUNDS

49. II. Types of Compounds IONIC COMPOUNDS COVALENT COMPOUNDS Formation: Ionic compounds form when one atom (usually a metal ) loses electrons to become a positively charged ion ( cation ), while another atom (usually a nonmetal ) gains those electrons to become a negatively charged ion ( anion ). The electrostatic attraction between the opposite charges results in the formation of an ionic bond.

50. II. Types of Compounds b. Examples: • Sodium Chloride (NaCl): Sodium (Na⁺) donates one electron to Chlorine (Cl⁻). • Magnesium Oxide (MgO): Magnesium (Mg²⁺) donates two electrons to Oxygen (O²⁻). IONIC COMPOUNDS COVALENT COMPOUNDS

51. II. Types of Compounds Properties: High Melting and Boiling Points : Due to the strong electrostatic forces between ions. Electrical Conductivity : Conducts electricity when molten or dissolved in water as ions are free to move. Solubility : Typically, soluble in water due to the polar nature of water molecules that interact with the ions. IONIC COMPOUNDS COVALENT COMPOUNDS

52. II. Types of Compounds IONIC COMPOUNDS COVALENT COMPOUNDS Formation: Covalent compounds form when two or more nonmetal atoms share pairs of electrons. This sharing allows each atom to attain a stable electron configuration , usually resembling that of the nearest noble gas.

53. II. Types of Compounds IONIC COMPOUNDS COVALENT COMPOUNDS b. Examples: Water (H₂O) : Oxygen shares one pair of electrons with each hydrogen atom. Methane (CH₄): Carbon shares one pair of electrons with each of the four hydrogen atoms. Carbon Dioxide (CO₂): Carbon shares two pairs of electrons with each oxygen atom

54. II. Types of Compounds IONIC COMPOUNDS COVALENT COMPOUNDS b. Properties: Lower Melting and Boiling Points : Compared to ionic compounds, because the forces holding molecules together (intermolecular forces) are generally weaker than ionic bonds. Nonconductive: generally do not conduct electricity in solid or molten form. Varied Solubility : Some covalent compounds are soluble in water (polar covalent) while others are not (nonpolar covalent).

55. II. Types of Compounds IONIC COMPOUNDS COVALENT COMPOUNDS

56. III. Polarity of Covalent Bonds Polarity refers to the distribution of electrical charge around atoms in a molecule . It affects properties like solubility, boiling and melting points, and reactivity

57. III. Polarity of Covalent Bonds Understanding Polarity Electronegativity: TWO TY PES OF COVALENT BOND: Nonpolar Covalent Bond Polar Covalent Bond:

58. III. Polarity of Covalent Bonds Understanding Polarity Electronegativity : The ability of an atom to attract shared electrons. The difference in electronegativity between two atoms in a bond determines the bond's polarity .

59. III. Polarity of Covalent Bonds Understanding Polarity TWO TYPES OF COVALENT BOND Nonpolar Covalent Bond : If the difference in electronegativity between the two atoms is small ( usually less than 0.5 ), the electrons are shared equally, resulting in a nonpolar bond.  

60. III. Polarity of Covalent Bonds Understanding Polarity TWO TYPES OF COVALENT BOND Polar Covalent Bond   If the electronegativity difference is significant ( between 0.5 and 1.7 ), the electrons are shared unequally, resulting in a polar bond where one atom has a partial negative charge (δ⁻) and the other a partial positive charge (δ⁺).

61. III. Polarity of Covalent Bonds Understanding Polarity TWO TYPES OF COVALENT BOND Polar Covalent Bond   If the electronegativity difference is significant ( between 0.5 and 1.7 ), the electrons are shared unequally, resulting in a polar bond where one atom has a partial negative charge (δ⁻) and the other a partial positive charge (δ⁺).

62. III. Polarity of Covalent Bonds Determining Polarity of Molecules Calculate the Electronegativity Difference (ΔEN): Subtract the electronegativity of one atom from the electronegativity of the other atom. Example: For HCl (Hydrogen Chloride), the electronegativity of H is 2.1, and Cl is 3.0. ΔEN = |3.0 - 2.1| = 0.9 Since 0.9 falls between 0.5 and 1.7 , HCl is a polar covalent molecule .

63. III. Polarity of Covalent Bonds Determining Polarity of Molecules Calculate the Electronegativity Difference (ΔEN): Calculate the electronegativity difference and determine if the following bonds are polar or nonpolar: H₂O (O = 3.44, H = 2.20) CH₄ (C = 2.55, H = 2.20) c. HCl (H = 2.20, Cl = 3.16) Example: For HCl (Hydrogen Chloride), the electronegativity of H is 2.1, and Cl is 3.0. ΔEN = |3.0 - 2.1| = 0.9 Since 0.9 falls between 0.5 and 1.7 , HCl is a polar covalent molecule .

64. I V . Chemical Formula Chemical formulas are symbolic representations that show the elements in a compound and their relative proportions Molecular Formula Displays the actual number of atoms of each element in a molecule. Example: Glucose (C₆H₁₂O₆) . Empirical Formula Shows the simplest whole-number ratio of atoms in a compound. Example: The empirical formula for glucose (C₆H₁₂O₆) is CH₂O.

65. VI. RULES IN NAMING COMPOUNDS NEXT MEETING

66. VI. RULES IN NAMING COMPOUNDS Naming Ionic Compounds (Metal + Nonmetal) 1. Simple Binary Ionic Compounds : Name the cation (metal) first, followed by the anion (nonmetal) with its ending changed to “-ide.” Example: NaCl = Sodium Chloride.

67. VI. RULES IN NAMING COMPOUNDS A. Naming Ionic Compounds (Metal + Nonmetal) 1. Simple Binary Ionic Compounds : MgF 2 K 2 O BeI 2 CaS AlBr 3

68. VI. RULES IN NAMING COMPOUNDS A. Naming Ionic Compounds (Metal + Nonmetal) 2. Transition Metals:

69. A. Naming Ionic Compounds (Metal + Nonmetal) 2. Transition Metals:

70. VI. RULES IN NAMING COMPOUNDS A. Naming Ionic Compounds (Metal + Nonmetal) 2. Transition Metals (Multivalent) Identify the METALS and verify if it forms more than one ion. Determine the ratio of ions and REVERSE the CROSSOVER. The positive and negative charges must BALANCE. Write the charge of the metal ion as ROMAN NUMERAL in brackets between the name of the metal and nonmetal. Do not forget that the ending of the NONMETAL is “IDE”.

71. VI. RULES IN NAMING COMPOUNDS A. Naming Ionic Compounds (Metal + Nonmetal) 2. Transition Metals: (Multivalent)

72. VI. RULES IN NAMING COMPOUNDS A. Naming Ionic Compounds (Metal + Nonmetal) 2. Transition Metals : (Multivalent) PbCl 2 FeF 2 Cu 3 N FeN

73. VI. RULES IN NAMING COMPOUNDS A. Naming Ionic Compounds (Metal + Nonmetal) 3. Polyatomic Ions: Retain the name of the polyatomic ion when naming the compound. Example: NaNO ₃ = Sodium Nitrate, where NO₃⁻ is the nitrate ion.

74. VI. RULES IN NAMING COMPOUNDS A. Naming Ionic Compounds (Metal + Nonmetal) 3. Polyatomic Ions: Retain the name of the polyatomic ion when naming the compound. Example: NaNO ₃ = Sodium Nitrate, where NO₃⁻ is the nitrate ion.

75. VI. RULES IN NAMING COMPOUNDS A. Naming Ionic Compounds (Metal + Nonmetal) 3. Polyatomic Ions: Retain the name of the polyatomic ion when naming the compound. CaSO 4 K 2 CO 3 Al(PO 4 ) NH 4 Cl Mg(NO 3 ) 2

76. VI. RULES IN NAMING COMPOUNDS Naming Covalent Compounds (Two Nonmetals) Prefixes: Indicate the number of atoms of each element:

77. VI. RULES IN NAMING COMPOUNDS Naming Covalent Compounds (Two Nonmetals) Example: CO = Carbon Monoxide CO₂ = Carbon Dioxide P₄O₁₀ = ???? Note: No “Mono-” on the First Element: If the first element has only one atom, the prefix “mono-” is typically omitted.

78. VI. RULES IN NAMING COMPOUNDS C. Naming Compounds with Polyatomic Ions The name of the compound is derived by combining the names of the cation and the anion. Example: NH₄Cl = Ammonium Chloride, where NH₄⁺ is the ammonium ion. CaCO ₃ = Calcium Carbonate, where CO₃²⁻ is the carbonate ion.

79. PRACTICE IDENTIFY IF IONIC OR COVLANENT COMPOUNDS: Iron Oxide Barium Chloride Carbon Dioxide Magnesium Oxide Nitrogen Tribromide

80. PRACTICE Write the name of each compound: SO 3 N 2 S PH 3 BF 3 P 2 Br 4

81. PRACTICE Write the name of each compound: NaBr CaO Li 2 S MgBr 2 Be(OH) 2

82. PRACTICE Write the chemical formula/ of each compound (Polyatomic): ______________ = NaNO 3 Calcium Carbonate = __________ _______________ = Li 2 SO 4 Beryllium Phosphide = _________ _________________ = Mg(OH) 2

83. PRACTICE PbBr 4 Co 2 S 3 Fe(MnO 4 ) 2 HgF 2