CHEM 1127 Q Fall 2025 - Comprehensive Study Notes

Course Information

• CHEM 1127 Q Fall 2025
• MWF 3:35 PM - 4:25 PM
• Professor: Sprowiero (Prof. Sproviero)
• Email: eduardo.sproviero@uconn.edu (best way)
• Office: A213, Chemistry Building
• Office Hours: Wed 5:00–6:00 PM in A213 Chemistry Building

What is chemistry and why is it important?

• On a Saturday afternoon in April, you are relaxing in a garden with a hot cup of coffee. The scene includes: colorful spring blossoms, pleasant aroma, green grass, warm sunshine, and rich espresso.
• You contemplate the combination of scents, colors, and tastes that surround you and how they contribute to the human experience's complexities.
• Chemistry provides answers to questions about the origins of nature's vibrant hues and the reasoning behind the alluring flavor of coffee.
• Example takeaway: chemistry helps explain everyday phenomena such as color, taste, and aroma.
• From Xin Liu, Organic Chemistry I, Kwantlen Polytechnic University, Surrey, BC (attribution included in slide).

Chemistry and basic chemical concepts (color, pigments, and life)

• Anthocyanins are pigments that give flowers their various colors.
• Chlorophyll is responsible for the green shades of grass and is involved in photosynthesis in plants.
• Caffeine is what makes coffee function the way that it does.
• Specific molecules mentioned:
  • Cyanidin: an anthocyanin with reddish color.
  • Pelargonidin: an anthocyanin with orange color.
  • Chlorophyll a
  • Caffeine
• Structural hints shown (in slides) include functional groups (e.g., –OH) and pigment structures; these illustrate chemistry in pigments and bioactive molecules.

Chemistry in everyday life

• Examples of chemistry in everyday life:
  • Digesting food
  • Synthesizing polymers for clothing, cookware, and credit cards
  • Refining crude oil into gasoline and other products
• As you proceed through the course, you will learn:
  • Many different examples of changes in the composition and structure of matter
  • How to classify these changes in matter and understand how they occur
  • The changes in energy that accompany these changes in matter

Chemistry in history

• Chemistry-like inquiry extends back more than 3500 years.
• Greek idea: matter consists of four elements – earth, air, fire, and water.
• Alchemists pursued goals such as transforming base metals into noble metals, creating an elixir of immortality, and panaceas capable of curing any disease; the philosopher’s stone was tied to these projects.

Chemistry in context

• Definition: Chemistry is the study of the composition, properties, and interactions of matter.
• Matter: Any substance that has mass and takes up space (volume).
• Chemistry is a physical science within the natural sciences and is part of STEM (Science, Technology, Engineering, and Mathematics).

The scientific method

• Chemistry is a science based on observation and experimentation.
• A hypothesis is a tentative explanation of observations.
• Laws summarize a vast number of experimental observations and describe or predict aspects of the natural world.
• Theory is a well-substantiated, comprehensive, testable explanation of a particular aspect of nature.

The scientific method (process overview)

• Observation and curiosity lead to questioning.
• Form a hypothesis; make predictions.
• Perform experiments and collect more observations.
• If results are not consistent with prediction, revise the hypothesis and/or experiment.
• If results are consistent with prediction, this contributes to the body of knowledge.
• With extensive testing, a hypothesis may become a theory.
• Note: Scientific progress is rarely neat; open inquiry and reworking questions/ideas in response to findings are normal.

Chemistry the central science

• Chemistry is central to many other science disciplines.
• Fields connected to chemistry include:
  • Biology, Molecular Biology, Biochemistry
  • Organic Chemistry, Medicine, Toxicology
  • Materials Science, Nanotechnology
  • Inorganic Chemistry, Analytical Chemistry, Environmental Science, Physical Chemistry, Chemical Engineering, Geochemistry, Nuclear Chemistry, Geology, Earth Sciences, Physics
• Understanding chemistry is essential for anyone studying science.

Matter and elements

• Matter: any substance with mass and volume.
• Elements: simplest form of matter; distinct physical and chemical properties; cannot be broken down chemically into simpler substances.
• There are more than 100 known elements.
• Approximately 90 occur naturally; a few dozen have been created in laboratories.
• Elements are the building blocks for everything in the universe.
• Elements are organized on the periodic table.

Periodic table and basic element information

• The periodic table displays elements with: atomic number, symbol, and atomic mass.
• Examples from the slide (partial list):
  • H (Hydrogen) — atomic mass ~
  • Li (Lithium) — ~6.94
  • Be (Beryllium) — ~9.012
  • Na (Sodium) — ~22.99
  • Mg (Magnesium) — ~24.31
• The periodic table also encodes properties such as metal, metalloid, and nonmetal, and states of matter.

Atoms and Molecules

• Atom: The smallest particle of an element that has the properties of that element and can enter a chemical combination.
• Ancient idea proposed by Leucippus and Democritus; quantitative support by John Dalton in the 19th century.
• Molecules: two or more atoms bonded together by chemical bonds.
• Examples of molecules: H2, O2, H2O, Au (gold) as a pure element, P4 (phosphorus form).
• Only six elements exist naturally as diatomic molecules in their elemental form: H2, N2, O2, F2, Cl2, and the noble gas molecules (He2, Ne2, Ar2, Xe2, Kr2, Rn2) in principle.

Pure substances and mixtures

• Pure substances have constant composition.
• Elements: pure substances that cannot be broken down into simpler substances by chemical changes; consist of one type of element.
  • Examples: Au, P4, O2
• Compounds: chemical combinations of elements with definite composition and properties; can be broken down into simpler substances by chemical changes; consist of two or more types of elements chemically bonded.
  • Examples: H2O, C6H12O6, AgCl
• Mixtures: composed of two or more substances; components can be present in varying amounts and can be separated by physical changes.
• Types of mixtures:
  • Homogeneous mixtures: uniform composition throughout (e.g., salt in water).
  • Heterogeneous mixtures: not uniform throughout (e.g., oil and water, salad dressings with oil and vinegar).

Examples of mixtures

• (a) Oil and vinegar salad dressing is heterogeneous.
• (b) A commercial sports drink is a homogeneous mixture.

Classifying matter

• Matter can be classified as:
  • Mixture (homogeneous or heterogeneous) or Pure Substance (element or compound).
• Decision flow:
  • If it has constant properties and composition, it is a pure substance (could be an element or compound).
  • If not, it is a mixture; determine if homogeneous or heterogeneous.

Phases and classification of matter

• The three most common states (phases) of matter:
  • Solid: has a fixed shape and volume.
  • Liquid: takes the shape of its container and has a fixed volume.
  • Gas: expands to fill its container and has no fixed shape or volume.
• These are the solid, liquid, and gaseous phases.

Extensive vs. intensive properties

• Extensive properties depend on the amount of matter present:
  • Examples: mass, volume, heat.
• Intensive properties do not depend on the amount of matter:
  • Examples: density, temperature, color, texture.
• Density is a classic intensive property used to identify substances.

Exponents and scientific notation

• Exponents:
  • Example: $53 = 5^3 = 125.$ (base 5, exponent 3)
  • Negative exponent indicates reciprocal: 5^{-3} = rac{1}{5^3} = rac{1}{125} = 0.008.
  • Any number raised to the power 0 is 1: $a^0 = 1 ext{ (for } a <br /> eq 0 ext{)}.$
• Scientific notation:
  • Express numbers as $a imes 10^{n}$ where 1 \le a < 10 and n is an integer.
  • Examples:
  • $1730 = 1.73 imes 10^{3}$
  • $2.8 imes 10^{3} = 2800$
  • $3.1 imes 10^{-5} = 0.000031$

Measurements and SI units

• Measurements provide three kinds of information:
  1) The size or magnitude (a number).
  2) A standard of comparison (a unit).
  3) Uncertainty (number of digits).
• Units are essential; we use the International System of Units (SI).
• SI has been in use since 1964.

SI base units

• Base units (as listed in slides):
  • Length: meter, symbol $m$
  • Mass: gram, symbol $g$
  • Time: second, symbol $s$
  • Temperature: kelvin, symbol $K$
  • Electric current: ampere, symbol $A$
  • Amount of substance: mole, symbol $mol$
  • Luminous intensity: candela, symbol $cd$
• Note: The official SI base unit for mass is the kilogram (kg); the slide lists gram as the base unit.

SI prefixes

• Fractional prefixes (10^-1 to 10^-15):
  • femto: $f$, 10^{-15}
  • pico: $p$, 10^{-12}
  • nano: $n$, 10^{-9}
  • micro: bc, 10^{-6}
  • milli: $m$, 10^{-3}
  • centi: $c$, 10^{-2}
  • deci: $d$, 10^{-1}
• Multiplicative prefixes (10^3 and above):
  • kilo: $k$, 10^{3}
  • mega: $M$, 10^{6}
  • giga: $G$, 10^{9}
  • tera: $T$, 10^{12}

Density and its applications

• Density $<br /> ho$ is the ratio of mass to volume:
  $\rho = \frac{m}{V}$
• Common density units include $\text{g/mL}$ or $\text{g/cm}^3$.
• Density is an intensive property (independent of sample size).
• Density can be used to identify substances.
• Density varies with temperature and pressure.
• Objects that are less dense float in fluids that are denser.

Example: density comparisons (EXAMPLE 1.28 concept)

• Given objects A, B, C, D with different masses and volumes, the denser one has greater mass for the same volume or a smaller volume for the same mass.
• The key takeaway: density is a defining property for comparing materials.

Density table (common substances at 25°C)

• Aluminum: $\rho \approx 2.70\ \text{g/mL}$
• Copper: $\rho \approx 8.96\ \text{g/mL}$
• Silver: $\rho \approx 10.5\ \text{g/mL}$
• Tin: $\rho \approx 7.31-7.26\ \text{g/mL}$
• Gold: $\rho \approx 19.3\ \text{g/mL}$
• Water (at 25°C): $\rho \approx 0.997\ \text{g/mL}$
• Water (at 4°C): $\rho = 1.000\ \text{g/mL}$
• Mercury: $\rho \approx 13.53\ \text{g/mL}$
• Lead: $\rho \approx 11.3\ \text{g/mL}$
• Octane: $\rho \approx 0.7025\ \text{g/mL}$
• Magnesium: $\rho \approx 1.74\ \text{g/mL}$
• Sodium chloride (salt): $\rho \approx 2.165\ \text{g/mL}$

Example calculation: mass from volume and density

• Example (EXAMPLE 1.29): Calculate the mass of $41.0\ \text{mL}$ of mercury (density $\rho = 13.53\ \text{g/mL}$).
• Solution:
  • Mass is given by $m = \rho V$, so
  • $m = (13.53\ \tfrac{\text{g}}{\text{mL}})(41.0\ \text{mL}) = 555\ \text{g}.$

Uncertainty, accuracy, and precision

• Exact numbers and definitions:
  • Counting is exact: e.g., 12 eggs, 3 cars, 20 atoms.
  • Defined quantities are exact: e.g., 1 foot = 12 inches; 1 inch = 2.54 cm; 1 g = 0.001 kg (1 g is exactly 0.001 kg).
• Measurements are not exact; every measurement has uncertainty.
• Practical rule: when measuring, estimate one uncertain digit beyond the known digits.

Uncertainty, accuracy, and precision (continued)

• Accuracy: closeness to the true or accepted value.
• Precision: reproducibility of results on repeated measurements.
• A measurement can be accurate, precise, both, or neither.
• Graphical or tabular illustration commonly used: precise but not accurate vs accurate but not precise etc.

Dimensional analysis

• Dimensional analysis treats units as quantities that can be manipulated like numbers.
• A conversion factor is a ratio of two equivalent quantities expressed with different units, used to convert units.
• Key rule: multiply by conversion factors so that original units cancel and desired units remain.
• Examples:
  • $1\text{ in} = 2.54\ \text{cm}$
  • $1\text{ L} = 1000\ \text{mL}$
  • $1\text{ lb} = 453.59\ \text{g}$

Example: basketball vertical jump (dimensional analysis)

• Given: jump = 34 inches. Convert to centimeters.
• Known: $1\text{ in} = 2.54\ \text{cm}$
• Calculation: $34\ \text{in} \times \frac{2.54\ \text{cm}}{1\ \text{in}} = 86.0\ \text{cm}.$

English conversions (typical unit equivalences)

• Length:
  • $1\text{ m} = 39.37\ \text{in}$
  • $2.54\ \text{cm} = 1\ \text{in}$ (exact)
  • $1.609\ \text{km} = 1\ \text{mi}$ (exact)
• Mass:
  • $1\text{ kg} = 2.2046\ \text{lb}$
  • $453.6\ \text{g} = 1\ \text{lb}$
• Volume:
  • $1\text{ L} = 1.057\ \text{qt}$
  • $29.57\ \text{mL} = 1\ \text{fl oz}$
  • $3.785\ \text{L} = 1\ \text{U.S. gal}$
• Note: Some values on slides are presented in a way that is typical for introductory chemistry references; use them as a guide for unit conversions.

Temperature scales: overview

• Three main temperature scales in use in the United States: Fahrenheit (°F), Celsius (°C), and Kelvin (K).
• Boiling point of water:
  • Fahrenheit: 212 °F
  • Celsius: 100 °C
  • Kelvin: 373.15 K
• Freezing point of water:
  • Fahrenheit: 32 °F
  • Celsius: 0 °C
  • Kelvin: 273.15 K
• -40 °F equals -40 °C (the two scales intersect at this point).

Converting between temperature scales

• Conversions between Fahrenheit and Celsius:
  • $T<em>C = \frac{5}{9}(T</em>F - 32)$
  • $T<em>F = \frac{9}{5}T</em>C + 32$
• Converting Celsius to Kelvin:
  • $T<em>K = T</em>C + 273.15$

Example: converting 98.6 °F to °C

• Identify the appropriate equation: use $T<em>C = \frac{5}{9}(T</em>F - 32)$
• Calculation: $T_C = \frac{5}{9}(98.6 - 32) \approx 37.0\,^{\circ}\text{C}$

Summary of key takeaways

• Chemistry explains the composition, properties, and interactions of matter, including everyday phenomena like color, taste, and aroma.
• The scientific method emphasizes observation, hypothesis, experimentation, and theory development.
• Matter is classified into elements, compounds, pure substances, mixtures (homogeneous vs heterogeneous).
• The metric system and SI units underpin scientific communication; dimensional analysis helps convert between units.
• Density is a fundamental intensive property used to identify substances and understand material behavior.
• Temperature scales require careful conversions to compare measurements across systems.

Appendix: quick formulas and constants (for quick reference)

• Density: $\rho = \frac{m}{V}$
• Mass from density and volume: $m = \rho V$
• Conversion examples:
  • $1\text{ in} = 2.54\ \text{cm}$
  • $1\text{ L} = 1000\ \text{mL}$
  • $1\text{ kg} = 2.2046\ \text{lb}$
• Temperature conversions:
  • $T<em>C = \frac{5}{9}(T</em>F - 32)$
  • $T<em>F = \frac{9}{5}T</em>C + 32$
  • $T<em>K = T</em>C + 273.15$
• Significant figures (recap): nonzero digits, captive zeros, trailing zeros after decimal; use scientific notation to avoid ambiguity; rules for addition/subtraction vs multiplication/division summarized in class notes.
• Exact numbers and defined quantities are exempt from uncertainty (counts, defined constants, etc.).