Chemistry Basics: Matter and Measurement

What is Matter?

Matter is anything that has mass and occupies space. Understanding matter is fundamental to chemistry, as it forms the basis for all substances and reactions.

• Definition: Matter includes solids, liquids, gases, and plasma.

• Classification: Matter can be classified as pure substances or mixtures.

• Pure Substances:

  • elements (cannot be broken down further)

  • compounds (composed of two or more elements chemically combined).

• Mixtures: Can be

  • homogeneous (uniform composition, e.g., salt water)

  • heterogeneous (non-uniform, e.g., salad).

• Example: Air is a homogeneous mixture; granite is a heterogeneous mixture.

Molecular Formula and Molecules

A molecular formula provides the types and numbers of atoms in a molecule, which is a group of atoms bonded together.

• Molecular Formula: Shows the exact number of each type of atom in a molecule (e.g., for water).

• Molecule: The smallest unit of a compound that retains its chemical properties.

• Example: is the molecular formula for carbon dioxide.

The Periodic Table

The periodic table organizes elements based on their properties and atomic structure.

• Period: A horizontal row; elements in a period have the same number of electron shells.

• Group/Family: A vertical column; elements in a group share similar chemical properties due to similar valence electron configurations.

• Important Elements for Human Nutrition: Includes elements like Iron (Fe), Fluorine (F), Iodine (I), and Lithium (Li).

• Iron: Essential for hemoglobin function in blood.

• Trace Metals: Often required for enzyme activity.

Physical States of Matter

Matter exists in different physical states, each with distinct properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume, no definite shape.

• Gas: No definite shape or volume.

• Plasma: Ionized gas, most common state in the universe (Additional info: not typically encountered in everyday life).

Physical and Chemical Changes

Physical changes alter the form of matter without changing its chemical identity, while chemical changes (reactions) produce new substances.

• Physical Change: Driven by energy (often heat); examples include melting, boiling, and dissolving.

• Chemical Reaction: Involves breaking and forming chemical bonds; produces new substances.

• Chemical Equation: Represents a reaction; reactants are transformed into products.

• Example:

International System of Units (SI) and Metric Prefixes

The SI system is the standard for scientific measurement, using base units and prefixes to indicate scale.

• Base Units: Meter (m), Liter (L), Kilogram (kg). Will also use cubic meters m^3

• Metric Prefixes: Indicate multiples or fractions of units (e.g., kilo-, milli-, micro-, nano-).

• Giga(G), Mega(M), Kilo(k), BASE UNIT, Deci(d), Centi(c), Milli(m), Micro(mc), Nano(n)

• Example: ; : 1000mg= 1g

• Usage: Choose appropriate prefixes for convenience (e.g., ng for small masses).

Unit Conversions and Dimensional Analysis

Unit conversions use conversion factors to change units without altering the quantity's value.

• Conversion Factor: A ratio used to convert from one unit to another.

• Dimensional Analysis: A method to ensure units cancel appropriately during calculations.

• Example:

Significant Figures and Scientific Notation

Significant figures reflect the precision of measurements; scientific notation expresses large or small numbers efficiently.

• Significant Figures: Digits that carry meaning in a measurement.

• Limitation: The precision of the measuring instrument limits significant figures.

• Sig Figs: if it has a decimal move left to right and start counting at the first non zero diget keep counting to the end. If NO decimal point move from right to left and do the same thing.

• Scientific notation: When numbers become smaller AKA moving the decimal to the left =positive exponent. When moving to the right, number get bigger negative exponent. If you are weird in the beginning you stay weird.

• Scientific Notation: Numbers written as .

• Example: becomes

Percentages and Fractions

Percentages are used to express proportions; converting between fractions and percentages is a common calculation.

• Conversion: Multiply fraction by 100 to get percentage.

• Example:

Mass, Weight, Volume, and Density

These are fundamental physical properties used in chemistry.

• Mass: Amount of matter in an object (measured in grams or kilograms).

• Weight: Force exerted by gravity on mass.

• Volume: Space occupied by matter (measured in liters, cubic centimeters).

• Density: Mass per unit volume;

• Example: ; of water weighs on Earth.

Specific Gravity and Refractometer

Specific gravity compares the density of a substance to water; refractometers measure this property.

• Specific Gravity:

• Importance: Used to assess purity and concentration; lipids are less dense than water, proteins are more dense.

• Refractometer: Instrument used to measure specific gravity.

Temperature and Its Scales

Temperature is a measure of thermal energy; conversions between scales are often required.

• Scales: Celsius (C), Fahrenheit (F), Kelvin (K).

• Conversion Formulas:

• Normal Human Body Temperature: F, C, $310$ K

• Hypothermia: Body temperature below normal.

• Hyperthermia: Body temperature above normal.

Energy, Calorie, and Joule

Energy is the capacity to do work; calories and joules are units of energy.

• Calorie (cal): Amount of energy needed to raise 1 g of water by 1C.

• Calorie (Cal): Used in nutrition;

• Joule (J): SI unit of energy;

• Potential Energy: Stored energy.

• Kinetic Energy: Energy of motion.

• Conservation of Energy: Energy cannot be created or destroyed.

Specific Heat Capacity

Specific heat is the amount of energy required to raise the temperature of 1 g of a substance by 1C.

• Formula:

• Water: Has a high specific heat capacity, which helps regulate body temperature.

Characteristics of States of Matter

Each state of matter has unique characteristics; transitions are driven by energy changes.

• Solid: Particles tightly packed, fixed shape.

• Liquid: Particles less tightly packed, flows.

• Gas: Particles far apart, fills container.

• Physical Changes: Driven by heat or energy input.

Accuracy and Precision in Measurement

Accuracy refers to how close a measurement is to the true value; precision is how reproducible measurements are.

• Accuracy: Correctness of a measurement.

• Precision: Consistency of repeated measurements.

• Example: A scale that gives the same value each time is precise; if that value is correct, it is also accurate.

U.S. Customary Units and Metric Conversions

Understanding both systems is important for practical applications in chemistry and medicine.

• Common Units: Pound (lb), quart (qt), ounce (oz), teaspoon (tsp), inch (in).

• Conversion: Use conversion factors to switch between systems (e.g., ).

Dosage Calculations and Drop Units

Calculating dosages and understanding drop units is essential in medical chemistry.

• Dosage: Amount of medication per patient, often based on body weight.

• Drop Units: indicates drops per milliliter.

• Percent Active Ingredient: Used to determine concentration in solutions.

• Percent of Adult Dose: Used for pediatric dosing.

Nutrition Labels and Macronutrients

Nutrition labels provide information on macronutrients and their percentages.

• Macronutrients: Carbohydrates, proteins, fats.

• Trans Fat: Important to monitor due to health risks.

• Calculation: Determine percentage of each macronutrient from total calories.

Table: Metric Prefixes (Main Purpose: Classification of Units)

Table: Common Physical States of Matter (Main Purpose: Comparison)

Table: Temperature Conversion Formulas (Main Purpose: Reference)

Table: Common U.S. Customary Units and Metric Equivalents (Main Purpose: Conversion)

Additional info: Some content was inferred for completeness, such as the inclusion of plasma as a state of matter and expanded explanations of trace metals and nutrition label calculations.