Physical and Chemical Properties, Measurement, and Significant Figures

Physical and Chemical Properties

Definition of Identity: In the context of chemistry, the term identity refers to the fundamental nature of a substance. Chemical changes involve a change in identity, whereas physical changes do not.
Definition of Properties: Properties are characteristics that enable us to distinguish one substance from another. Each substance possesses a unique set of properties (e.g., a lump of coal versus a lump of silver).
Physical Properties: - A physical property is a characteristic of matter that is not associated with a change in the chemical composition or identity of a substance. - Examples include density, color, hardness, melting point, boiling point, and electrical conductivity. - State of Matter Changes: Phase changes like evaporation, condensation, boiling, and melting are physical properties. - The Water Scenario: If water is boiled on a stove to make spaghetti, it turns from liquid to gas. If it condensates on a lid, it turns back to liquid. Through both changes, the identity remains water ( $H_2O$ ). - Melted Butter Scenario: Melting butter for a recipe changes its state from solid to liquid, but its identity as butter remains unchanged.
Chemical Properties: - A chemical property involves the change of one type of matter into another, signifying a fundamental change in identity. - Common chemical properties include flammability, toxicity, acidity, reactivity, and heat of combustion. - Combustion/Flammability Example: If a wooden stick is set on fire, it burns and becomes ash. The ash cannot be considered wood; it is a different substance (wood ash) resulting from a chemical change. - Digestion Example: Food consumed is dissolved by stomach acid, undergoing chemical reactions to transform into energy for the body. - Rusting (Oxidation) Example: Iron ( $Fe$ ) reacts with oxygen to form iron oxide ( $Fe_2O_3$ ), commonly known as rust. This changes the chemical identity. In contrast, Chromium ( $Cr$ ) is often used for vehicle trim because it does not rust.
Indicators of Chemical Change: - Color Change: A common indicator that a chemical reaction has occurred. Specific examples include: - The formation of brown gas during certain reactions. - Cooking meat: The oxidation of iron in myoglobin causes chicken or hamburger meat to change from red to brown. - Ripening Bananas: Bananas change flavor and color (green to yellow to brown) via a chemical process where the fruit effectively "self-feeds" to mature. This process involves the production of ethylene gas. - Burning a Match: Striking a match results in a chemical change; once burned, the match cannot be struck again because its properties have fundamentally changed.

Extensive and Intensive Properties

Extensive Properties: - These properties depend on the amount of matter present. - Examples: Volume and mass. For instance, the amount of heat required to warm a large frying pan is greater than for a small one because it has more mass.
Intensive Properties: - These properties do not depend on the amount of matter present. - Examples: Density and temperature. - Density of Water: Water always has a density of approximately $1.0\,g/cm^3$ regardless of whether you have a small bottle or 50 liters of it.

The NFPA Hazard Diamond

The National Fire Protection Agency (NFPA) uses a hazard diamond to relay safe handling information regarding chemical and physical properties.
Structure of the Diamond: - Red (Top): Fire hazard (Flammability). It indicates the flash point, which is the temperature at which a substance will combust. - Blue (Left): Health hazard. A rating of 3 indicates an extremely dangerous substance that should not be handled with bare hands. - Yellow (Right): Reactivity. This indicates how likely a substance is to undergo detonation or violent chemical changes (e.g., releasing significant heat). - White (Bottom): Specific hazards. This indicates unique properties, such as "Use No Water," which is common for reactive metals like Sodium ( $Na$ ).

The Periodic Table as a Tool for Predicting Properties

The periodic table is an essential tool for predicting the properties of elements and compounds.
Historical Context: - Dmitri Mendeleev: Often considered the "Father of the Periodic Table." - Henry Moseley: Independently arrived at similar conclusions regarding the arrangement of elements.
Organization: - Elements were originally categorized by their known properties at the time (e.g., those that react violently with water or stable gases). - Mendeleev initially organized elements by increasing atomic mass and was able to predict elements that were not yet discovered based on empty spots in the property trends. - Groups: Vertical columns on the table are called groups. Elements within a group share similar chemical and physical properties and follow specific trends.

Fundamentals of Measurement

Quantitative vs. Qualitative: Science prioritizes quantitative measurements (numerical values) over qualitative surveys (humanities style).
The Three Parts of a Measurement: 1. A Number: The value of the quantity. 2. A Unit: Provides the standard by which the number is measured. Without a unit, a number like "13.25" is meaningless. Units distinguish between vastly different quantities (e.g., $10\,grams$ vs. $10\,tons$ ). 3. Uncertainty: Indicates how accurate or precise the measurement is based on the instrument used. More decimal places correlate to higher certainty.
Measurement Systems: - English System: Used primarily in the US; includes units like the foot, pound, and gallon. - Metric System (SI): Used by the rest of the world and in science; includes the meter, liter, and kilogram. Note that "Metric" and "SI" (International System of Units) have been used interchangeably since 1964.

The Metric (SI) System and Prefixes

SI Base Units: - Length: Meter ( $m$ ). - Mass: Kilogram ( $kg$ ). - Time: Second ( $s$ ). - Temperature: Kelvin ( $K$ ). - Amount of Substance: Mole ( $mol$ ). - Electric Current: Ampere ( $A$ ). - Luminous Intensity: Candela ( $cd$ ).
Metric Prefixes: Used to express fractional or multiple SI units by factors of 10 (most often factors of $10^3$ ). - Smaller prefixes: deci ( $10^{-1}$ ), centi ( $10^{-2}$ ), milli ( $10^{-3}$ ), micro ( $10^{-6}$ ), nano ( $10^{-9}$ ), pico ( $10^{-12}$ ), femto ( $10^{-15}$ ). - Larger prefixes: kilo ( $10^3$ ), mega ( $10^6$ ), giga ( $10^9$ ), tera ( $10^{12}$ ).

Mass, Length, Time, and Temperature

Length: The meter was originally defined as one ten-millionth of the distance from the North Pole to the Equator. It is now defined as the distance light travels in a vacuum in $\frac{1}{299,792,458}$ of a second. A meter is roughly 3 inches longer than a yard.
Relationship between Inches and Centimeters: $1\,inch = 2.54\,centimeters$ .
Mass: The kilogram was originally the mass of one liter of water. It is standardized by a platinum-iridium alloy cylinder kept in France (a prototype is also kept at the NIST in Maryland). One kilogram is roughly $2.2\,lbs$ .
Temperature: - Kelvin is the SI unit. It is written as $K$ , not $\text{degrees } K$ . - The magnitude (increment) of 1 Kelvin is equal to 1 degree Celsius ( $^∘C$ ). - Freezing/Boiling points of water: - Celsius: Freezes at $0^∘C$ , boils at $100^∘C$ . - Kelvin: Freezes at $273.15\,K$ , boils at $373.15\,K$ . - Absolute Zero: $0\,K$ is $-273.15^∘C$ . - Typical Temperatures: Room temperature is roughly $25^∘C$ ( $298\,K$ ); body temperature is roughly $37^∘C$ ( $310\,K$ ).
Time: The SI unit is the second ( $s$ ). It is incorrect to use "sec."

Derived Units: Volume and Density

Derived Units: Units that combine two or more base units.
Volume: The measure of space occupied by an object. It is length cubed ( $L \times W \times H$ ). - The SI unit is the cubic meter ( $m^3$ ), but chemistry often uses the liter ( $L$ ) and milliliter ( $mL$ ). - Crucial Equivalence: $1\,cm^3 = 1\,mL$ .
Density: A ratio of mass to volume. - Formula: $D = \frac{m}{V}$ . - Units: Common units are $g/mL$ or $g/cm^3$ . - Density can be used as a conversion factor. For example, knowing water is $1.0\,g/mL$ , $10\,grams$ of water will occupy $10\,mL$ of volume.

Uncertainty and Significant Figures

Exact Numbers: These have no uncertainty. - Counting: You cannot have half a chicken; counting 10 chickens is an exact measurement. - Defined Quantities/Conversion Factors: $12\,inches = 1\,foot$ and $1\,inch = 2.54\,cm$ are exact values.
Measured Numbers: These always contain uncertainty due to practical limitations of instruments.
Reading Measurements: - Always record all certain digits plus one estimated (uncertain) digit. - Meniscus: When reading a graduated cylinder, water curves downward. Always read from the lowest point of the meniscus. - If a cylinder has markings for every $1\,mL$ , and the water is between $21$ and $22$ , the reading might be reported as $21.6\,mL$ . If the cylinder has half-tick marks ( $0.5\,mL$ ), and the water is between $21.5$ and $22$ , the reading might be $21.55\,mL$ .
Significant Figure Rules: 1. Non-zeros: All non-zero digits are significant (e.g., $13.52$ has 4 sig figs). 2. Interior Zeros: Zeros between non-zeros are always significant (e.g., $10,052$ has 5 sig figs). 3. Leading Zeros: Zeros to the left of the first non-zero digit are never significant (e.g., $0.00345$ has 3 sig figs). 4. Trailing Zeros: - They are significant if there is a decimal point (e.g., $120.0$ has 4 sig figs). - They are not significant if there is no decimal point (e.g., $1,200$ has 2 sig figs).