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Types of Matter

Types of Matter. Matter must. Have mass Have volume (take up space). Matter exists in three phases. Solids (s)- fixed shape & volume Liquid (l)- fixed volume, takes the shape of the container Gas (g)- takes both the volume and the shape of the container.

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Types of Matter

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  1. Types of Matter

  2. Matter must • Have mass • Have volume (take up space)

  3. Matter exists in three phases • Solids (s)- fixed shape & volume • Liquid (l)- fixed volume, takes the shape of the container • Gas (g)- takes both the volume and the shape of the container

  4. Matter is divided into 2 categories. • Pure Substances • Fixed composition • Unique set of properties • Mixtures • Composed of two more substances physically mixed

  5. Pure Substances • Elements-type of matter than cannot be broken down into two or more pure substances • Elements can be divided into many categories • Metals • Nonmetals • Metalloids • Families (Periodic Table) • Compounds-more than one element • Have new properties after chemically combined

  6. Mixures • Homogeneous-composition is the same throughout • A solution is a liquid homogeneous mixture made of solvents & solutes. • Most solvents are liquids; however it can be a gas • Heterogeneous- does not have a uniform composition

  7. Separation Techiques We are going to look at 3 types. There are more. • Filtration • Used for a heterogeneous solid-liquid mixture • Distillation • Homogeneous solid-liquid mixture • Also liquid-liquid mixtures in which one liquid can be evaporated • Chromatography • Liquid-liquid mixture • Liquid-gas mixture • Gas-gas mixture

  8. Significant Figures • Uncertainty of at least one unit in the last digit • 2.00 mL = 3 SF • 2.0 mL = 2 SF • 2 mL = 1 SF • 0.002 mL= 1 SF

  9. Rules • Zeros appearing between nonzero digits are significant • Zeros appearing in front of all nonzero digits are NOT significant. • Zeros at the end of a number and to the right of a decimal point are significant. • A decimal point placed after zeros indicates that they are significant. • Trailing zeros without decimals are questionable • 75000 may be 2-5 SF • 7.5 x 104 = 2 SF • 7.50 x 104= 3 SF • Assume not significant in the book therefore 75000= 2SF

  10. Rules • Multiplying & Dividing • The # of SF in the result is the same as quantity of smallest # of SF • Adding & Subtracting • The number of decimal places in the result is the same with the smallest decimal • Rounding • If less than 5… leave last digit unchanged • If greater than 5…add one to the last digit • If equal to 5…round even • Round 14.575 to 2SF= 14

  11. Rules • Multiple Operations • Carry out all steps with complete number of digits • Go back to find the # of SF at each step • Round at the end • ALWAYS SHOW WORK AND USE YOUR UNITS!!!

  12. How many Sig Figs? • 0.00800 in • 52.000 nm • 800 ns • 4.30 x 104 kg • 5060 g

  13. Solve the problem with correct SF. Box in your answer. • 2.505 x 0.0920 x 451.08= • 0.0810 + 7.168 + 1.50 = • 5.20/8.973= ***When putting in your calculator remember to use () to make calculator do order of operations correctly.

  14. Conversions • When you make a conversion, choose the factor that cancels out the initial unit. • The actual conversion factor does not count on significant figures • Use the number of significant figures in the initial measurement given • Remember what % means mathematically. When given a % use the unit shown in the problem. • 52% mL = 52 mL/ 100 mL • 0.07% L = 0.07 L/ 100

  15. Examples • 3 hr to sec • 5.27 Mg to lb (Mg= mega grams) • 0.53 mi to m • 55.25 mi to km • 11.6 mL to in3

  16. Properties of Substances • Every pure substance has its own unique set of properties • Chemist use these properties for identification • 2 ways of grouping these properties

  17. Intensive & Extensive Properties • Intensive Properties (do not depend on amount) ***chemist use these to identify • These are a few examples • Density • Melting pt • Boiling pt • Extensive Properties (depend on amount ) • Mass • Volume

  18. Physical & Chemical Properites • Physical- do not change the substance • Theses are just a few examples • Density • Melting pt • Boiling pt • Solubility • Color • Chemical-when a new substance is formed • Reactivity • Flammability

  19. Density • Denisty= mass/volume • Mass= grams • Volume= cm3, mL, L • Be sure to show all work and units!!!!!

  20. Examples • The density of Al is 2.70 g/mL. What is the volume of 8.21 grams? • The density of Hg is 13.5 g/mL. How many grams are in 5.0 mL?

  21. Examples • A sample of metal in a small weighing dish had a mass of 59.61g. The dish had a mass of 0.58 g. When the metal was added to the water, the water level rose 9 mL. • What is the mass of the metal? • What is the density of the metal?

  22. Examples • A sample of metal with a mass of 10.06g was placed in a flask with a volume of 65.0 mL. To fill the flask, 35.7 g Hg (density= 13.5 g/mL) must be added to the metal. What is the density of the metal?

  23. Moles to grams to particles • Atomic Mass= mass on the periodic table • Avogadro’s Number • Symbol NA • 6.022 x 1023 • It represents the number of atoms of an element in a samples whose mass in grams is equal to the atomic mass of that element • 6.022 x 1023 H atoms= 1.008 g H • 6.022 x 1023 N atoms = 14.001 g N

  24. Examples • Find the mass of a N atom • Find the number of N atoms in 7.00 grams.

  25. Moles and Molar Mass • Moles • 1 mole = Avogadro’s Number • 1 mole = 6.022 x 1023 • Molar Mass • Units are g/mol • Same as formula mass

  26. Examples • 13 g of caffeine, C4H5N2O • Convert to moles • Convert from moles to atoms • Convert from atoms to Number of Carbon atoms

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