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Fundamental Units and States of Matter

This review covers fundamental units of measurement in the metric system and the different states of matter. It also discusses physical and chemical properties, changes, and the structure of atoms, elements, and compounds.

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Fundamental Units and States of Matter

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  1. Final Review 268 Slides – I would suggest listening and trying to find what topics you are weak in then studying and putting that info into your notes for tomorrow’s final. 100 multiple choice questions (open notes) covers all of the chapters.

  2. 7 Fundamental Units (also referred to as SI “System International” Units) Masskilogramkg Lengthmeter m Time second s TemperatureKelvin K Count, quantity molemol Electric current ampere A Luminous intensity candela cd

  3. Examples of Non-SI Units: Temperature °C Pressure atm Volume liter Energy calorie Force psi

  4. Examples of Derived Units combinations of SI basic units Area meter squared m2 Volume meter cubed m3 • The metric system makes these fundamental units larger or smaller by using Prefixes

  5. Prefixes MegaM 1,000,000 Kilo k 1,000 Hectah100 Deca da 10 UNIT g, m, L, etc. 1 Deci d 1/10 Centi c 1/100 or .01 Milli m 1/1000 or .001 Micro μ 1/1,000,000 Nano n 1/ 1,000,000,000

  6. Mass and weight are not identical • Mass is how much stuff is in a sample while weight is dependent on the gravity of the planet. • Typical lab unit for mass? ____Gram (g)_____ • Typical lab unit for length? __centimeter (cm)___ • Typical lab unit for liquid volume? milliliter (mL)_ • Typical lab unit for solid volume? cubic centimeters (cm3) • NOTE: mL3 is not a measurement!

  7. Density Ratio of the amount of mass in a given volume Density = mass volume Typical units for density: SOLID ____g/cm3_______ LIQUID ____g/mL_______ NOTE: Useful tip: _1 cm3 = 1 mL__

  8. States of Matter Solid definite shape and volume Ex. Desk, Gold Liquid definite volume; shape of container Ex. Water at Room Temp Gas takes shape and volume of container Ex. Air, Oxygen Plasma exists only at high temperatures Ex. Sun, Stars

  9. Physical property Characteristic of a substance that can be altered without changing the identity of the substance • Ex. odor, color, volume, state at room temperature, density, melting point, boiling point, attraction to a magnet

  10. Chemical property Ability of a substance to change to a different substance • Ex. flammability, corrosion

  11. Physical change Change that does not affect the identity of a substance Ex. Rip Paper Ex. Melting Silver Ex. Salt Dissolves in Water

  12. Chemical change Change that alters the identity of a substance Ex. Burn Paper Ex. Silver Tarnishes Ex. Electrolysis of Water Ex. Adding Silver Nitrate to Sodium Chloride 4 indicators that a chemical reaction occurred: Evolution of a Gas Spontaneous Color Change Spontaneous Temp Change Formation of a Precipitant

  13. The Atom Atom- the basic unit of matter Element- substance made up of only one type of atom that cannot be broken down into simpler substances by ordinary means Ex. iron, oxygen, sodium Very rare to find isolated elements; they are usually combined with other elements!

  14. Compounds Two or more elements chemically combined in a fixed proportion that can be broken down into simpler substances by chemical means Ex. pure water, sugar, sodium chloride NOTE: The properties of the individual elements are often different than those of the compound (ex. hydrogen and oxygen are gases, H2O is a liquid)

  15. Substances Pure substance- substance that always has the same composition Ex. elements and compounds Mixture- a physical blend of two or more pure substances Ex. soda, coffee, air, drinking water Mixtures can be separated into their pure substances by physical means Helpful Hint: How to decide if something is a compound or a mixture

  16. Types of mixtures Homogeneous mixture- mixture with no visibly different parts; uniform throughout; also called a solution (NOT only liquids) Ex. salt water (H2O and NaCl) , brass (Cu and Zn), air (N2, O2, others) Heterogeneous mixture- mixture with visibly different parts Ex. sand and water, glass of ice water

  17. Elements, Atoms and Ions Orbit Atomic Structure General concept of an atom: Atom- The smallest part of an element that retains the chemical identity of that element Planetary Model of The Atom Nucleus Quantum Mechanical Model of The Atom Orbital

  18. Subatomic Particles

  19. Atomic number The number of protons Symbol: _Z_ Ex. Mg is atomic number 12. It has 12 protons Ex. Tungsten is atomic number _74_. It has _74_ protons.

  20. Isotopes Atoms of the same element (the same number of protons) that have different numbers of neutrons. This means they have the same atomic number but different mass numbers (the next concept).

  21. Mass number the number of protons plus the number of neutrons (NOT a number value found on the periodic table) Symbol: _A_ Ex. Na with a mass number of 23 (can be written as Na-23) Ex. Na with a mass number of 24 (can be written as __Na-24__) ***Common Misunderstanding: Note that both of these are isotopes of sodium!

  22. Three Isotopes of Hydrogen hydrogen-1 hydrogen-2 hydrogen-3 Recall, mass number = P + N 1 = 1 + N 2 = 1 + N 3 = 1 + N N = 0 N = 1 N = 2 Again, all three are isotopes of hydrogen. Hydrogen-2 is NOT considered the “normal” atom.

  23. Isotope Notation

  24. Symbol for hydrogen-3

  25. Losing Electrons Sodium atom (+11) + (-11) = zero net charge Sodium loses 1 electron (+11) + (-10) = +1 charge Cation- positive ion; formed when electrons are lost; usually metals Symbol for cations – Xcharge   Ex. Na+1 Ex. Mg+2 Naming cations- use name of the parent atom plus the word “ion” Ex. Na+1 = Sodium ion Ex. Mg+2 = Magnesium ion

  26. Gaining Electrons: Chlorine atom (+17) + (-17) = zero net charge Chlorine atom gains 1 electron (+17) + (-18) = -1 charge Anion- negative ion; formed when electrons are gained; nonmetals Symbol for anions – Xcharge Ex. Cl-1 Ex. S-2 Naming anions- take root word, add ending “-ide”, and the word “ion”   Ex. Cl-1 = Chloride ion Ex. S-2 = Sulfide ion

  27. Three categories of elements Metals- most elements; left side and middle of periodic table Properties of metals: • conduct heat and electricity • malleable– hammered into sheets • ductile – drawn into a wire • lustrous– shiny • most are solid at room temperature (except Hg) • most are silver in color (except Cu and Au) EXCEPTION: HYDROGEN IS NOT A METAL (it is a nonmetal)

  28. Three categories of elements Metalloids/ semi-metals- mixture of metallic and nonmetallic properties; zigzag line (a.k.a. stairstep line) Ex. boron- B antimony- Sbsilicon- Si tellurium- Te germanium- Gepolonium- Po arsenic- As astatine- At • EXCEPTION: ALUMINUM IS NOT A METALLOID (it is a metal)

  29. Diatomics Most elements do not occur in nature by themselves. Diatomic molecule- element, when found alone in nature (not combined with other elements) is always found in pairs Mnemonics: HONClBrIFor HONClFIBror “Magic Seven”   The seven diatomic elements are: H2,O2,N2,Cl2,Br2,I2,F2

  30. Examples Ex. CO2 = carbon dioxide Ex. NO = nitrogen monoxide Ex. N2O5 = dinitrogenpentoxide

  31. Common Names • Common Names of Molecular Compounds- these are not named according to the system above and you WILL have to MEMORIZE these! H2O = Water NH3 = Ammonia CH4 = Methane (actually is the base of a whole different form of nomenclature that we are not going to cover in this class)

  32. Extra Practice:

  33. You MUST memorize the names, formulas and charges of these six polyatomic ions: +1 ions NH4+ ammonium ion -1 ions OH- hydroxide ion NO2- nitrite ion NO3- nitrate ion HCO3- hydrogen carbonate ion (bicarbonate) C2H3O2- acetate ion CN- cyanide ion -2 ions SO3-2 sulfite ion SO4-2 sulfate ion CO3-2 carbonate ion -3 ions PO4-3 phosphate ion

  34. Formula Writing Naming acids will be covered later in the year. For now, you only need to know the formulas: hydrochloric acid = HCl sulfuric acid = H2SO4 nitric acid = HNO3 acetic acid = HC2H3O2

  35. Hydrates Solid compounds that contain water in their crystalline structure Ex. CuSO4 5H2O = copper (II) sulfate pentahydrate They are easily named by naming the ionic compound part and then using a prefix for the number of water molecules that are present followed by the word “hydrate” Ex. BaCl2 2H2O = barium chloride dihydrate

  36. Chemical Reactions • A chemical change is really a chemical reaction. Four signs a chemical change has occurred: • color change • formation of a precipitant • evolution of a gas • spontaneous change in temperature

  37. Chemical reaction • Chemical reaction- bonds (attraction between atoms) are broken, atoms are rearranged and new bonds form • Chemical equation- used to represent a chemical reaction • Reactants- substances present before a reaction • Products- substances present after a reaction

  38. Another Example • Methane gas reacts with oxygen gas to produce carbon dioxide gas and water vapor. Add the physical state symbols next to each substance CH4(g) + O2(g)  CO2(g) + H2O(g)

  39. Practice Example: Potassium metal dissolves in liquid water to form hydrogen gas and aqueous potassium hydroxide. Reactants: K + H2O Products: H2 + KOH K(s) + H2O(l) H2(g) + KOH(aq)

  40. Practice Example: Aqueous potassium iodide and aqueous lead (II) nitrate reacts to form aqueous potassium nitrate and a precipitate of lead (II) iodide. Reactants: KI + Pb(NO3)2 Products: KNO3 + PbI2 KI(aq) + Pb(NO3)2(aq) KNO3(aq) + PbI2(s)

  41. Now you try a problem like that described above from start to finish: Ex: Liquid ethanol (C2H5OH) reacts with oxygen gas to form carbon dioxide and water vapor. Answer: C2H5OH(l) + 3O2(g) 2CO2 (g) + 3H2O(g)

  42. Classifying Reactions • Directions: Predict the products and then balance the following equations. You do not need to include the physical states. Synthesis (direct combination or composition)- a compound is formed from simpler substances Ex. 2Na + Cl2 → 2NaCl A + B  AB

  43. Classifying Reactions Decomposition- a compound is broken down into simpler substances Ex. 2H2O → 2H2 + O2 AB  A + B

  44. Classifying Reactions Single replacement (displacement)- cations or anions are exchanged Ex. Mg + CuSO4 → MgSO4 + Cu (Mg and Cu+2 switch) (metal replaces a metal) A + BC  AC + B SHOULD - Need to check the reactivity series for metals to see if a single metal replacement reaction occurs.

  45. Classifying Reactions Double replacement reaction- reaction in which cations and anions are exchanged (like replaces like) Ex. CaCO3 + 2HCl → CaCl2 + H2CO3 AB + CD  AD + CB

  46. Classifying Reactions Combustion- a rapid reaction involving oxygen that produces heat (flame) Hint: Sometimes these are difficult to balance! Ex. CH4 + 2O2 → CO2 + 2H2O (products are always the same) CxHx + O2 → CO2 + H2O CH4is a hydrocarbon – a compound containing hydrogen and carbon

  47. Classifying Reactions Acid/Base Neutralization Reaction: is a chemical reaction in which an acid and a base react with each other. Usually produces water and a salt. Hint: Sometimes these are difficult to balance! Ex. KOH + HCl→ KCl+ H2O (Salt – any ionically bound material) Mg(OH)2+ 2HNO3→ Mg(NO3)2+ 2H2O

  48. Ethene gas (C2H4) is burned in oxygen gas to produce carbon dioxide gas and water vapor.

  49. Examples Directions for turning large numbers into standard scientific notation: Ex. 1000 = 1 x 103 Ex. 125 = 1.25 x 102 Ex. 17,000 = 1.7 x 104 Ex. 93,000,000 = 9.3 x 107

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