Chemistry chapter 1

• Chemistry is a subdivision of physical science that focuses on what happens when the electron cloud of one substance encounters the electron cloud of another
• The Scientific Method involves observations, formulating hypotheses, and performing experiments to validate hypotheses
• When two atoms collide, electrons may be transferred, leading to ionization and the formation of an ionic bond
• Chemistry is considered the central science, with a covalent bond having a length of approximately 10^-10 meters
• When two hydrogen atoms collide, electrons may be shared, forming a covalent bond
• Matter is anything that occupies space and has mass. A substance is a form of matter with a definite composition and distinct properties. Chemistry is the study of matter and the changes it undergoes
• The three states of matter are solid (fixed shape and volume), liquid (varying shape but fixed volume), and gas (no fixed shape or volume)
• A mixture is a combination of two or more substances where they retain their distinct identities. It can be homogeneous (same composition throughout) or heterogeneous (non-uniform composition)
• Physical properties are unique characteristics of a material, while chemical properties relate to how it reacts with other materials
• A burning candle is a chemical change as it involves the breaking and formation of new chemical bonds
• Physical changes involve a change of state or dissolving, while chemical changes result in the formation of different substances through the rearrangement of atoms and the formation of new chemical bonds
• Most real-world processes involve both physical and chemical changes. For example, when a candle burns, it undergoes both physical (melting) and chemical (burning) changes
• Exponential notation is used in chemistry to simplify writing large or small numbers, with scientific notation being the preferred method
• Positive exponents indicate how many places to the left the decimal has been moved
• Numbers with positive exponents are large
• Negative exponents indicate how many places to the right the decimal has been moved
• Numbers with negative exponents are small
• Scientific notation involves moving the decimal point to express numbers in a more concise form
• Calculations with scientific notation:

  • For addition/subtraction, write each quantity with the same exponent and then add or subtract the numbers
  • For multiplication, multiply the numbers and add the exponents
  • For division, divide the numbers and subtract the exponents
• Significant figures are important in scientific calculations to indicate the precision of a measurement
• Significant figures do not include uncertain digits or place-holding zeroes
• Determining significant figures:

  • If no decimal is present, trailing zeroes are not significant
  • If a decimal is present, leading zeroes are not significant
• Significant Figures:

  • 0.0003007 has 4 significant figures
  • 15.00 has 4 significant figures
  • 1200 has 2 significant figures
• Scientific Notation:

  • 2.3 x 10^-3 written as a non-exponential number is 0.0023
  • 56,100,000,000 in scientific notation is 5.61 x 10^10
• When converting numbers to scientific notation, the exponent is determined by the number of places the decimal point is moved
• When performing calculations with significant figures, the result should have no more significant figures than the least precise value used in the calculation
• Significant Figures:

  • 0.0030 L has 2 significant figures
  • 0.1044 g has 4 significant figures
  • 53,069 mL has 5 significant figures
  • 0.00004715 m = 4.715 x 10^-5 m has 4 significant figures
  • 57,600 s = 5.76 x 10^4 s has 3 significant figures
  • 0.0000007160 cm^3 = 7.160 x 10^-7 cm^3 has 4 significant figures
• Unit prefixes:

  • Giga- (G) = 10^9
  • Mega- (M) = 10^6
  • Kilo- (k) = 10^3
  • Deci- (d) = 10^-1
  • Centi- (c) = 10^-2
  • Milli- (m) = 10^-3
  • Micro- (μ) = 10^-6
  • Nano- (n) = 10^-9
  • Pico- (p) = 10^-12
• How many grams are in 350 milligrams (mg)?

  • 350 mg is equal to 0.350 g or 3.50 x 10^-1 g
  • Replace the "m" for "milli-" with the exponent from the prefix table, which means 10^-3
• Temperature Scales:

  • Metric: Celsius (or centigrade)
  • SI: Kelvin
  • English: Fahrenheit
• Celsius vs. Kelvin:

  • Both scales have the same size of degree
  • Celsius goes from 0 to 100, while Kelvin goes from 273 to 373
• Fahrenheit Scale:

  • Has 180 steps between freezing and boiling points of water
  • 1°C = 1.8°F
  • The scales differ in their starting points: 32°F = 0°C
• Temperature Conversions:

  • K = °C + 273
  • °C = K - 273
• If the temperature outside is 263 K, the equivalent in °C is:

  • °C = 263 - 273
  • °C = -10 (Feels like winter!)
• Celsius/Fahrenheit Conversions:

  • °F = (1.8 x °C) + 32
  • Adjustments for the difference in degree size and zero point offset are made
• To convert Fahrenheit to Celsius:

  • Subtract 32 from the Fahrenheit temperature
  • Multiply the result by 5/9
  • The result is the temperature in Celsius
• To convert Celsius to Fahrenheit:

  • Multiply the Celsius temperature by 1.8
  • Add 32 to the result
  • The final temperature is in Fahrenheit
• Formula to convert Celsius to Fahrenheit:

  oF = (1.8 x oC) + 32
• Formula to convert Fahrenheit to Celsius:

  oC = (oF - 32) / 1.8
• Unit Conversions:

  Length:

  • 1 km = 0.6214 mile
  • 1 m = 1.094 yards
  • 1 m = 39.37 inches
  • 1 cm = 0.3937 inch

  Volume:

  • 1 cubic meter (m^3) = 35.31 ft^3
  • 1 dm^3 = 0.2642 gal
  • 1 dm^3 = 1.057 qt
  • 1 cm^3 = 0.03381 fluid ounce

  Mass:

  • 1 kg = 2.205 lb
  • 1 g = 0.03527 ounce (oz)
  • 1 lb = 0.4536 kg
  • 1 oz = 28.35 g