Chapter 2-3: Measurements, Units, and Matter & Energy (VOCABULARY Flashcards)

Units and Measurement Systems

• Chemistry text uses Metric (cgs/Si-prefixed) and SI units throughout Chapter 2–3; key base units include Volume (L, mL, m³), Length (m), Mass (g, kg), Temperature (°C, K), Time (s).

• Table highlights: Volume = L (liters), Length = m (meters), Mass = g → kg, Temperature = °C → K, Time = s.

• Useful conversions (as given):

  • $1\ \text{L} = 1000\ \text{mL}$

  • $1\ \text{L} = 1.06\ \text{qt}$

  • $1\ \text{mL} = 1\ \text{cc}$

  • $1\ \text{qt} = 2\ \text{pints}$

  • $1\ \text{m} = 100\ \text{cm}$

  • $1\ \text{m} = 39.4\ \text{in}$

  • $1\ \text{m} = 1.09\ \text{yd}$

  • $2.54\ \text{cm} = 1\ \text{in}$

  • $1\ \text{kg} = 1000\ \text{g}$

  • $1\ \text{kg} = 2.20\ \text{lb}$

  • $454\ \text{g} = 1\ \text{lb}$

• Temperature notes:

  • Celsius scale: water freezes at $0^ ext{\circ}C$ and boils at $100^ ext{\circ}C$.

  • Kelvin scale (K) is SI; 0 K is absolute zero; relationship: $T<em>K = T</em>C + 273.15$.

• General approach to measurement:

  • A measuring tool yields measured numbers; the last digit is the estimated digit.

  • Estimated digit represents uncertainty; report it as the final digit (e.g., 4.5 cm).

Volume, Length, and Mass Relationships

• Volume relations:

  • $1\ \text{L} = 1000\ \text{mL}$

  • $1\ \text{L} = 1.06\ \text{qt}$

  • $1\ \text{mL} = 1\ \text{cm}^3 = 1\ \text{cc}$

• Length relations:

  • $1\ \text{m} = 100\ \text{cm}$

  • $1\ \text{m} = 39.4\ \text{in}$

  • $1\ \text{m} = 1.09\ \text{yd}$

  • $2.54\ \text{cm} = 1\ \text{in}$

• Mass relations:

  • $1\ \text{kg} = 1000\ \text{g}$

  • $1\ \text{kg} = 2.20\ \text{lb}$

  • $454\ \text{g} = 1\ \text{lb}$

Temperature and Its Scales

• Temperature measures how hot or cold an object feels; expressed on different scales:

  • Celsius scale (°C)

  • Kelvin scale (K) in SI

• Key points:

  • Water freezes at $0^ ext{\circ}C$ and boils at $100^ ext{\circ}C$.

  • Kelvin scale begins at absolute zero: $0\ \text{K}$ corresponds to $-273.15^ ext{\circ}C$.

• Common formulas:

  • $T<em>K = T</em>C + 273.15$

  • $T<em>C = T</em>K - 273.15$

Measured Numbers and Significant Figures

• Measured numbers come from a measuring tool; they have estimated digits.

• Significant figures (SFs) include all digits that are known with certainty plus the last estimated digit.

• SF rules (summary):

  • All nonzero digits are significant.

  • Zeros may or may not be significant:

  • Leading zeros are not significant.

  • Confined zeros (zeros between nonzero digits) are significant.

  • Trailing zeros: significant only if a decimal point is present.

• Examples (illustrative):

  • 4.5 g has $2\ \text{SF}$

  • 122.35 m has $5\ \text{SF}$

  • 205 °C has $3\ \text{SF}$

  • L has $2\ \text{SF}$

  • 16.00 mL has $4\ \text{SF}$

• Zeros in scientific notation keep only significant zeros:

  • g → $3\ \text{SF}$

  • 5.00 × 10^2 g → $3\ \text{SF}$

  • 0.30 cm → $2\ \text{SF}$

  • 3.0 × 10^1 cm → $2\ \text{SF}$

Exact Numbers

• Exact numbers are counted values or defined relations; they have unlimited SF and do not limit the SF in calculations.

• Examples: 8 donuts, 1 L = 1000 mL, 1 foot = 12 inches, 1 km = 1000 m, 1 lb = 16 oz.

Significant Figures in Calculations

• Rounding rules (general):

  • If the first dropped digit is 4 or less, drop it and all following digits.

  • If the first dropped digit is 5 or greater, increase the last retained digit by 1.

• Calculations with SFs:

  • Multiplication/Division: the result has as many SFs as the factor with the fewest SFs.

  • Example (illustrative): if 2.8 (2 SF) × 67.40 (4 SF) × 34.8 (3 SF) → final with 2 SF; product ≈ 6.7 × 10^3.

  • Addition/Subtraction: the result should have the same number of decimal places as the quantity with the fewest decimal places.

  • Example: 2.012 (3 decimals) + 61.09 (2 decimals) + 3.0 (1 decimal) → 66.1 (tenths place).

• Practice rounding examples: 3.145 g rounded to 3 SF and 2 SF.

Prefixes (Metric/SI)

• Prefixes that increase size:

  • Tera (T, 10^12), Giga (G, 10^9), Mega (M, 10^6), Kilo (k, 10^3)

  • Examples: 1 km = 10^3 m, 1 Gm = 10^6 m, 1 Tg = 10^12 g (in context, units matter).

• Prefixes that decrease size:

  • Deci (d, 10^-1), Centi (c, 10^-2), Milli (m, 10^-3), Micro (µ, 10^-6), Nano (n, 10^-9), Pico (p, 10^-12)

• Special note: in medicine, micro (µ) is sometimes written as mc for safety reasons (e.g., mcg).

Measuring Volume and Volume Equivalents

• The cubic centimeter (cc) is the volume of a cube 1 cm on each side.

• 1 cc = 1 mL (and 1 L = 1000 mL).

Converting Between Units and Equalities

• Equalities: different units may describe the same quantity (metric-to-metric, US customary, or mixed).

• Common equalities (examples):

  • $1\ \text{km} = 1000\ \text{m}$

  • $1\ \text{m} = 1000\ \text{mm}$

  • $1\ \text{m} = 100\ \text{cm}$

  • $1\ \text{cm} = 10\ \text{mm}$

  • $1\ \text{L} = 1000\ \text{mL}$

  • $1\ \text{m}^3 = 1000\ \text{L}$

  • $1\ \text{kg} = 1000\ \text{g}$

  • $1\ \text{lb} = 16\ \text{oz}$

  • $1\ \text{gal} = 4\ \text{qt}$

• Use conversion factors as fractions to cancel units (e.g., 60 min / 1 h or 1 h / 60 min).

Problem Solving Using Conversion Factors

• Steps:

  • Identify given unit and needed unit.

  • Plan: write a chain of conversions to reach the needed unit.

  • Include one or more conversion factors that cancel units and provide the needed unit.

  • Perform calculation and check units.

• Example concept: converting body weight from pounds to kilograms using $1\ \text{kg} = 2.20\ \text{lb}$.

Density and Specific Gravity

• Density: mass per volume; $\rho = \frac{m}{V}$.

• Density relates to whether objects float or sink in water: density > 1 g/mL sinks; density < 1 g/mL floats.

• Specific gravity (sp. gr): a relative density; $\text{sp.gr} = \frac{\rho}{\rho_{\text{water}}} = \frac{\rho}{1.00\ \text{g/mL}}$.

Classification of Matter

• Matter is anything with mass and occupies space.

• Classifications by composition:

  • Pure substances: fixed composition; includes elements (one type of atom) and compounds (two or more elements in a fixed proportion).

  • Mixtures: two or more substances physically combined; can be separated by physical methods; variable composition.

• Mixtures can be:

  • Homogeneous: uniform composition throughout (e.g., Brass: copper + zinc).

  • Heterogeneous: not uniform; distinct phases visible (e.g., copper metal and water).

Physical States of Matter

• Solids: definite shape and volume; particles close together; strong interactions; e.g., ice, iron.

• Liquids: takes shape of container; definite volume; particles less rigid than solids.

• Gases: fill container; no fixed volume or shape; particles far apart; essentially no interactions.

• Table contrasts shape, volume, particle arrangement, and movement across solid, liquid, gas.

Physical vs Chemical Properties and Changes

• Physical properties: color, shape, odor, density, melting point, etc.; observed without changing identity.

• Chemical properties: relate to a substance's ability to form new substances (e.g., burn, rust, tarnish).

• Physical changes: change in state or form without changing composition (e.g., melting, grinding).

• Chemical changes: transform substances into new substances with new properties (e.g., burning, rusting).

• Examples and identifications (from text): classify given descriptions as physical vs chemical; changes vs properties.

Temperature Scales and Conversions (In-Text Examples)

• Fahrenheit, Celsius, and Kelvin scales; have reference points for water’s freezing/boiling.

• Conversions:

  • From Celsius to Fahrenheit: $F = 1.8 C + 32$

  • From Fahrenheit to Celsius: $C = (F - 32) / 1.8$ or $C = \frac{5}{9}(F - 32)$

  • From Celsius to Kelvin: $K = C + 273.15$

  • From Kelvin to Celsius: $C = K - 273.15$

Heat, Energy, and Food Energy Values

• Heat (thermic energy) is related to particle motion; faster motion = more heat.

• SI unit of energy/work: joule (J); $1\ \text{kJ} = 1000\ \text{J}$.

• Calorie definitions:

  • 1 cal raises 1 g water by 1°C; 1 kcal = 1000 cal; 1 cal = 4.184 J (exact for the definition).

  • 1 cal = 4.184 J; thus 1 kcal = 4184 J.

• Food energy values (typical macronutrient energy per gram):

  • Carbohydrate: $4\ \text{kcal/g} = 17\ \text{kJ/g}$

  • Fat: $9\ \text{kcal/g} = 38\ \text{kJ/g}$

  • Protein: $4\ \text{kcal/g} = 17\ \text{kJ/g}$

• Example: how many calories come from a cup of milk with given grams of carb/fat/protein; use kcal/g values above.

Specific Heat and Phase Changes

• Specific heat capacity (SH) varies by substance; units: $\text{J g}^{-1} \ \text{K}^{-1}$ or $\text{cal g}^{-1} \ \text{K}^{-1}$.

• Common reference: for water, $c<em>{H</em>2O} = 4.184\ \text{J g}^{-1}\ \text{K}^{-1} = 1.00\ \text{cal g}^{-1}\ \text{K}^{-1}$.

• Heat equation: $q = m c \Delta T$ where m is mass, c is specific heat, and \Delta T is change in temperature.

• Example: calculating heat gained or lost by a mass during a temperature change.

Phase Changes and Heats of Transformation

• Melting/Fusion: solid to liquid at the melting point; freezing is the reverse.

• Heat of fusion (ΔH_fus): energy to convert 1 g of solid to liquid at the melting point.

  • Example: for ice, ΔH_fus ≈ 80 cal/g = 334 J/g at 0 °C.

• Vaporization/Condensation: liquid to gas (boiling) and gas to liquid (condensation).

• Heat of Vaporization (ΔH_vap): energy to convert 1 g of liquid to gas at boiling point; for water, ≈ 540 cal/g = 2260 J/g at 100 °C.

• Sublimation/Deposition: solid to gas and gas to solid without liquid phase (e.g., dry ice).

Heating and Cooling Curves

• Heating curve: diagonal lines indicate temperature rise; horizontal plateaus indicate phase changes.

• Cooling curve: similar interpretation for temperature decrease.

• Plateaus occur at phase transition temperatures (fusion/vaporation points).

Worked Example Highlights

• Density measurement: given mass and displaced water volume, compute density from mass/volume.

• Specific gravity example: density relative to water (water density assumed $1.00\ \text{g/mL}$).

• Temperature problems: convert between °C, °F, and K using the relationships above.

• Problem solving with conversion factors: build calculation using cancellation of units; ensure units cancel and final unit matches target.

Practice and Learning Checks (Overview)

• Concept checks include identifying measured vs exact numbers, determining SFs, performing rounding, and applying SF rules in multiplication/division and addition/subtraction.

• Typical problems involve converting units, applying densities, and using heat equations for phase changes.

Quick Reference Tables (Key Points)

• Density: $\rho = \frac{m}{V}$

• Heat: $q = m c \Delta T$ and for phase changes:$q_ ext{fus} = m \Delta H_ ext{fus}, \quad q_ ext{vap} = m \Delta H_ ext{vap}$

• Temperature conversions:$F = 1.8T + 32,\quad C = \frac{F-32}{1.8},\quad K = C + 273.15$

• Energy values (food): Carbohydrate = $4\ \text{kcal/g}$, Fat = $9\ \text{kcal/g}$, Protein = $4\ \text{kcal/g}$; in kJ: $4\ \text{kcal/g} \approx 17\ \text{kJ/g}, \ 9\ \text{kcal/g} \approx 38\ \text{kJ/g}$

• 1 kcal = 4.184 kJ and 1 cal = 4.184 J

• 1 L = 1000 mL; 1 mL = 1 cm³; 1 m³ = 1000 L

• Prefix examples: kilo- (10^3), centi- (10^-2), milli- (10^-3), micro- (10^-6), nano- (10^-9), pico- (10^-12)

• Exact numbers do not limit SFs and do not affect calculation precision

• Plateaus on heating curves correspond to phase changes (fusion, vaporization)

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