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Chemistry
 

Chemistry

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  • Jawbreaker example
  • Valence – the ones in the outer shell. Same number of valence electrons similar properties.
  • Table salt Different properties of the elements they are made from.
  • Sodium donates electron to chlorine to help fill its outer orbital.
  • Sea water, salt water, air, bronze (copper and tin), brass (copper and zinc)

Chemistry Chemistry Presentation Transcript

  • Earth/Environmental Science Chemistry Chapter 3 P Squires 2005-2006
  • The study of matter is central to the study of chemistry.
    • Q. What is matter?
    • A. Matter is anything that …
    • … has mass,
    • … and takes up space.
  • Do things you cannot see have mass?
    • What about air?
  • What is density?
    • Density is the ratio of the mass of an object to the volume of the object.
    Typical units of density are grams per milliliter, g/mL D = m V
  • Develop a method to measure the density of a piece of metal.
    • What equipment will you need?
    • What data should you take?
    • How will you analyze the data?
  • Devise a method to measure the density of a liquid.
    • What laboratory equipment will you need?
    • What kind of data should you take?
    • How will you analyze the data?
  • Temperature is a measure of “how hot or cold” something is. How do we measure temperature?
  • Temperature Scales K C F 100 0 -273 373 273 0 212 32 -462 Boiling point of water Freezing point of water (Melting point of ice) Absolute zero
  • Physical States
    • Solid
    • Liquid
    • Gas
    • Plasma
    What are some characteristics of each?
  • Physical and Chemical Properties
    • Physical properties only involve changes in state or appearance.
    • Chemical changes and chemical properties always produce new substances.
  • Some physical properties …
    • Color
    • Density
    • Melting point
    • Boiling point
    • Solubility
    Conductivity Odor Hardness Malleability Ductility
  • Elements
    • A substance that cannot be broken down into simpler substances by physical or chemical means.
    • Ninety-two elements occur naturally.
    • Other elements have been produced in labs.
  • Elements
    • Symbols are used to identify the different elements.
  • Most abundant elements in universe.
    • Hydrogen
    • Helium
  • Elements
    • Do MiniLab on page 55.
    • ONLY DO FIRST TWO COLUMNS
    • (We will do the others later in class.)
  • Atoms
    • The smallest particles of an element that has all the characteristics of that element.
  • Dalton’s Atomic Theory
    • Atoms are tiny, discrete particles
    • Atoms are indestructible
    • Atoms of the same element have the same mass and properties
    • Atoms combine in simple whole-number ratios
    • Atoms in different ratios produce different compounds.
  • Nucleus
    • The center of the atom that is made up of protons and neutrons .
    • Protons are tiny particles that have mass and a positive (+) electrical charge.
    • Neutrons are particles with about the same mass but no charge.
  • Atomic and Mass Number
    • Atomic number - the number of protons in the nucleus of an atom
    • Mass number - the number of protons added to the number of neutrons
  • Electrons
    • The number of electrons in a neutral atom equals the number of protons.
  • Electrons
    • Surround the nucleus of the atom.
    • Have very little mass.
    • Have a negative (-) charge.
  • Electrons
    • Exist in energy levels.
  • Hydrogen atom nucleus Discrete energy levels for electrons electron
  • Valence Electrons
    • The electrons in the outer shell.
    • Want to have 8 valence electrons for a full shell.
  • Periodic Table 1 2 3 4 5 6 7 IA IIA IIIA …………….IIA IIIA………… VIIIA
  • Periodic Table
    • Metals – Hard and shiny; good conductors of heat and electricity
    1 2 3 4 5 6 7
  • Periodic Table
    • Non-metals: Dull and brittle; poor conductors of heat and electricity.
    1 2 3 4 5 6 7
  • Group 1A: Alkali Metals
    • React with water easily to lose a valence electron.
    • Form a +1 ion
    • Soft
    • Found in nature combined with other elements
    • Good conductors of heat and electricity
  • Group 2A: Alkaline Earth Metals
    • Forms compounds with Oxygen (Oxides)
    • React with water
    • Shiny solids that are heavier than Alkalis
    • React with Oxygen or other metals found in the earth
    • Lose valance electrons to form +2 ions
  • Group 7A: Halogens
    • Form compounds with almost all metals
    • Salts (NaCl)
    • Have different Physical properties
    • Very reactive and found combined in nature
    • Tend to share or gain 1 e- to form a –1 ion
  • Group 8A: Noble Gases
    • Colorless
    • Unreactive
    • Last to be discovered
    • Rarely react (1 st time was in a lab in 1962)
  • Transition Metals
    • Metals in the middle.
    • Share properties of metals, just not as strong.
    1 2 3 4 5 6 7
  • Lanthanides and Actinides 1 2 3 4 5 6 7 Man made elements. Would make table too big if inserted where they go.
  • Metalloids
    • Found along the stair step on periodic table.
    • Have properties of metals and nonmetals.
  • Isotopes
    • Number of neutrons vary.
    • Example:
        • Chlorine atom
        • 17 protons + 18 neutrons
        • 17 protons + 20 neutrons
  • Activity
    • Finish mini lab on page 55.
  • Isotopes
    • Isotopes are what causes radiation and are part of the study of nuclear chemistry.
  • Compounds
    • Substance that is composed of atoms of two or more different elements that are chemically combined.
  • Chemical Bonds
    • Forces that hold elements together in a compound.
  • What are the two types of bonds? Ionic bonds - transfer electrons to get more stable arrangements. Covalent bonds - share electrons to get more stable arrangements. What happens to the electrons?
  • Ionic Bond Na Cl +
  • Na + Makes Na + ions and Cl - ions + - An ionic compound: NaCl Cl
  • Covalent Bond
  • Bonding
  • Molecule
    • A compound of two or more atoms held together by covalent bonds.
  • Chemical Formula
    • H 2 O
    • water
  • Polar Molecules
    • Do not share electrons equally.
  • Ions
    • An atom that either gains or loses electrons.
    • Is a charged particle. (+) or (-)
    • If has less than 4 valence electrons…tends to lose electrons
    • If has more than 4 valence…tends to gain electrons.
  • Ionic Bonding
    • Attractive forces between two atoms.
    • Example: NaCl
  • Activity
    • Do Problem Solving Lab page 63.
  • Metallic Bonds
    • An array of positive metal ions surrounded by a sea of mobile electrons.
  • Molecular Geometry
    • The structure of the molecule when formed.
    • Based on attraction of electrons.
  • VSEPR Theory V alence S hell E lectron P air R epulsion
  • Linear Molecular Geometry 180 degrees Two electron pairs Linear electron-pair geometry
  • Trigonal Planar Molecular Geometry Three electron pairs Trigonal planar electron-pair geometry 120 degrees
  • Tetrahedral Molecular Geometry Four electron pairs Tetrahedral electron-pair geometry 109.5 degrees
  • Making Molecular Models
    • Do making molecular models worksheet.
    • Carefully make the models and draw them in color on the worksheet.
    • Only complete 1 st and 2 nd column
  • Naming Chemical Elements
  • Naming Chemical Compounds
  • Chemical Reactions S(s) + O 2 (g)  SO 3 (g) 3 2 2 BCl 3 (g)  B(s) + Cl 2 (g) 2 2 3
  • Balancing Chemical Reactions Reactants: Zn + I 2 Product: Zn I 2
    • The chemical equation for this reaction ,
    • C + O 2  CO 2
  • Chemical Equations
    • 4 Al(s) + 3 O 2 (g) ---> 2 Al 2 O 3 (s)
    • This equation means
    • 4 Al atoms + 3 O 2 molecules ---produces--->
    • 2 molecules of Al 2 O 3
    • AND/OR
    • 4 moles of Al + 3 moles of O 2 ---produces--->
    • 2 moles of Al 2 O 3
  • Balancing Equations
    • ___ Al(s) + ___ Br 2 (l) ---> ___ Al 2 Br 6 (s)
    2 3
  • Balancing Equations
    • ____C 3 H 8 (g) + _____ O 2 (g) ----> _____CO 2 (g) + _____ H 2 O(g)
  • Balancing Equations ____B 4 H 10 (g) + _____ O 2 (g) ----> ___ B 2 O 3 (g) + _____ H 2 O(g)
  • Mixtures
    • A combination of two or more components that are easily recognizable.
    • What are some examples of mixtures?
  • Solutions
    • Part cannot be easily distinguishable, but have all of original properties.
    • What are some examples of solutions?
  • Acids and Bases
    • Acids taste sour, and …
    • bases taste bitter.
    Acids are the H + ion. Bases are the OH - ion.
  • Acids and bases in the body. Acids and bases are found throughout the body. The most obvious is stomach acid. Also amino acids and proteins to name a few.
  • Stomach Acid …
    • Digestion in the stomach involves HCl and enzymes that break food apart.
    Sometimes the contents of the stomach move up into the esophagus and heartburn occurs because the esophagus is not protected by mucus.
  • … and Antacids
    • Antacids relieve heartburn.
    Antacids contain hydroxides and carbonates. Antacids contain bases and neutralize acids. Because of hydrolysis, carbonate compounds can act as bases. CO 3 2- + HOH  HCO 3 - + OH -
  • … and Antacids Gaviscon magnesium carbonate – MgCO 3 aluminum hydroxide – Al(OH) 3 Mylanta calcium carbonate – CaCO 3 magnesium hydroxide – Mg(OH) 2 TUMS calcium carbonate – CaCO 3
  • Acids and Bases
    • Acids and Bases neutralize each other.
    • Their strength is based off the pH scale.
  • The pH Scale…
    • pH measures the amount of acidity in a solution.
    Acidic Basic Neutral 0 ------------ 7 ------------ 14
  • Common Food Acidities
  • States of Matter
    • Solids
    • Liquids
    • Gases
    • Plasma
  • Crystalline Structure
    • Structures where particles are arranged in regular geometric patterns.
  • Geo Lab
    • Perform Geo Lab on page 70-71.
    • Use 75 mL of tap water.
    • Use 27 g of salt (NaCl)
    • 25 ml in one beaker for refrigerator, 25 ml heated in beaker, and 25 ml to sit in window.
    • Work in groups of 4 (two lab groups).
    • Make sure to label each beaker with your group name and block number.
  • Solids
    • Have definite shape and volume.
  • Liquids
    • Take the shape of their container.
    • Have definite volume.
  • Gases
    • Have no definite shape or volume.
    • Expand or contract to fill space available.
  • Plasma
    • Hot, highly ionized electrically conducting gases.
    • Example: lightening, neon sign
  • Changes of State
  • Evaporation
    • Changing from a liquid to a gas.
    • Aka: Vaporization
  • Sublimation
    • Changing from a solid to a gas without becoming a liquid.
    • Example: dry ice
  • Condensation
    • Change from a gas to a liquid.
  • Law of Conservation of Matter
    • Matter can neither be gained or lost only changed in form.
  • Law of Conservation of Energy
    • Energy can neither be lost or gained only changed in form.
  • Chapter 3 Assessment
    • What particles make up the nucleus of an atom?
    • a. protons only
    • b. neutrons only
    • c. neutrons and electrons
    • d. protons and neutrons
  • Chapter 3 Assessment
    • Which of these make up the nucleus of an atom?
    • a. number of protons
    • b. number of neutrons
    • c. neutrons and protons
    • d. protons and electrons
  • Chapter 3 Assessment
    • Which is the average of the mass numbers of an elements isotopes?
    • a. atomic number
    • b. energy levels
    • c. atomic mass
    • d. valence electrons
  • Chapter 3 Assessment
    • What is the most abundant element on the Earth’s crust?
    • a. hydrogen
    • b. uranium
    • c. silicon
    • d. oxygen
  • Chapter 3 Assessment
    • What kind of ions characterize an acid?
    • a. Hydroxide ions (OH - )
    • b. Hydrogen ions (H + )
    • c. Oxygen ions (O -2 )
    • d. Negative ions (-)
  • Chapter 3 Assessment
    • How many valence electrons do beryllium atoms (atomic number 4) have?
    • Why?
  • Chapter 3 Assessment
    • Why don’t gases such as neon and argon combine chemically with other elements?