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Chemical bonding

Chemical bonding. Objectives: Can you identify ionic and covalent bonds? Can you name chemical compounds?. Properties of Ionic Compounds. They are crystalline solids at room temperature (like salt) They have high melting points

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Chemical bonding

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  1. Chemical bonding Objectives: Can you identify ionic and covalent bonds? Can you name chemical compounds?

  2. Properties of Ionic Compounds • They are crystalline solids at room temperature (like salt) • They have high melting points • They conduct electricity when they are dissolved in water or when they are in molten stat e(heated up so high they turn into liquid) • They are made up of metals and nonmetals

  3. Ionic compounds • What holds ionic compounds together? • The difference in charges between the ions. • Opposites attract. • How are ions formed? • When electrons are transferred from one element to another. • The atom becomes an ion because the number of protons and electrons are no longer equal creating charges.

  4. Conductivity • Ionic compounds are non-conductors but they are conductors when in aqueous solution. • What does that mean? • When ionic solids are melted or dissolved, they break into ions. • The electrically charged ions are then free to move. • Their motion conducts electricity • Strong vs Weak Electrolytes • What are electrolytes? • Conductivity videos

  5. Ionic Bonding • If an element gains electrons, the charge is negative • If an element loses electrons, the charge is positive.

  6. Review periodic table • Each group has a specific charge. • Groups are columns. Look at your periodic table and label the charges associated with the group.

  7. Ionic bonds • Involves electrostatic attraction between opposite charges of atoms. • Cations • Positively charged atoms. • Anions • Negatively charged atoms. • Ends in -ide • Cations and Anions come together to form ionic compounds • Examples: • Sodium + Chlorine  Sodium Chloride • Na++ Cl- NaCl

  8. Ionic Bonds continued • What if the charges are different? • Example: Lithium (1+) and Oxygen(2-) • Criss-cross the numbers and bring them down. • Why? • Because the charges are different in ratio 1:2, criss-crossing will cancel the charges out to make a stable compound. • Transition elements with different charges have their charges in parenthesis like Iron (II) Fe2+ and Iron (III) Fe3+ , or Copper (II) Cu2+ and Copper (III) Cu3+

  9. Molecular/Covalent bonding • Instead of using charges to bond, covalent bonds share electrons. • Also called a molecular compound. • They can have unequal amount of electrons • This causes polarity • One side has more electronegativity • Polarity • Depends on shape of the molecule and orientation of bonds • Nonpolar substances have an equal amount of sharing and there is no pulling • Electronegativity – The tendency of an atom to attract electrons toward itself

  10. Characteristics of covalent compounds • Gases, liquids, and brittle solids (like Carbon) at room temperature • Low melting points • Do not conduct electricity when dissolved in water • Made up of nonmetals • Covalent compounds video

  11. Properties of Molecular compounds Appearance of pure molecular compounds • Most molecular compounds are liquids or gases at room temperature • Some are solids and have crystalline structure and are brittle. • Some are solids and have a powder-like appearance and are soft.

  12. Chemical Bonds • Why do atoms bond? (phet stimulation) • Stability • Interactions/attractions • Chemical bond – electrical attractions that hold atoms together and lower the energy state of the system (they become more stable)

  13. Covalent bonding nomenclature • If a compound has two or more nonmetals, they are written with prefixes of the number of that atom. • Example: • CO2 Carbon Dioxide • N2H4DinitrogenTetrahydride • 1 – Mono (for the second element) 6 - Hexa • 2 – Di 7 - Hepta • 3 – Tri 8 – Octa • 4 – Tetra 9 - Nona • 5 – Penta 10 – Deca • If the next word after the prefix starts with the vowel the “a” in the prefix is dropped. Example: Hexafluoride vs. Hexoxide, not Hexaoxide.

  14. Polyatomic ions • Groups of atoms that are covalently bonded together and have a charge on the whole group. • They function like any other single atom ion and combine with any charged ion. • Examples: • Sulfate SO42- • Carbonate CO32- • Nitrate NO3- • Ammonium NH4+

  15. Polyatomic ions continued • Most polyatomic ions have different amount of hydrogens oxygens and they are named according to how much more or how much less hydrogen or oxygen the polyatomic ion has. • Chlorate vs. Chlorite vs. Perchlorate vs. Hypochlorite (there will never be a hypochlorate or perchlorite) • -ate = the median amount of oxygen atoms • -ite = one less oxygen than –ate • Hypo= one less oygen than –te • Per = one more oxygen than -ate • Hypo-, -Ite, -Ate, Per- • Sufate vs. Sulfate. • Nitrate vs. Nitrite

  16. Polyatomic ion practice

  17. Acids Nomenclature • Elements that are combined with Hydrogen become binary acids. • HCl, HBr, HF • Nomenclature: Hydro+(element)-ic + acid. • Hydrochloric acid • Hydrofluoric acid • What is Bromine + Hydrogen? • Polyatomic ions that combine with Hydrogen become ternary acids • HSO4, HClO3 • Nomenclature: Ends with –ous or –ic + acid • -ate-ic, -ite-ous • Sulfate SO42- with Hydrogen  H2SO4 (Sulfuric acid) • Nitrite NO3- with H+  HNO3 (Nitrous acid) • What is Nitrate + Hydrogen?

  18. Acids nomenclature practice

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