Chemistry For Today: General, Organic, and Biochemistry - Chapters 1, 3, and 4

Matter, Mass, and Weight

• Matter: Defined as anything that possesses mass and occupies physical space.
• Mass: A measurement quantifying the amount of matter present in an object. Unlike weight, mass is independent of the object's location; an object possesses the identical mass on Earth as it does on the moon.
• Weight: A measurement of the gravitational force exerted on an object. This value depends entirely on the object's location. For example, an object weighing approximately $16\,pounds$ on Earth would weigh only about $2.7\,pounds$ on the moon.

Physical and Chemical Properties and Changes

• Physical Properties: Characteristics of matter that can be observed or measured without attempting to alter the chemical composition of the substance being observed. Common examples include color and size.
• Chemical Properties: The properties that matter displays when attempts are made to transform the substance into entirely new substances. An example is the ability of paper to burn.
• Physical Changes: Changes that occur without altering the chemical composition of the substance. For instance, reducing the size of a sheet of paper by cutting off a piece is a physical change.
• Chemical Changes: Transformations in matter that lead to a change in chemical composition. A primary example is the burning of magnesium metal.

Molecules and Atoms

• Molecule: The smallest particle of a pure substance that retains the properties of that substance and is capable of a stable, independent existence. This represents the ultimate limit of physical subdivision for a pure substance.
• Atoms: These represent the limit of chemical subdivision for matter.
• Classification of Molecules:   - Diatomic molecules: Composed of exactly two atoms.   - Triatomic molecules: Composed of exactly three atoms.   - Polyatomic molecules: Composed of more than three atoms.   - Homoatomic molecules: Molecules containing only one kind of atom (e.g., Oxygen, $O_2$).   - Heteroatomic molecules: Molecules containing two or more different kinds of atoms (e.g., Carbon monoxide $CO$, Carbon dioxide $CO_2$).

Classification of Matter

• Mixtures: Matter in which the proportions of components may vary, and properties vary with the composition. Mixtures can be physically separated into two or more pure substances.
• Pure Substance: Matter characterized by a constant composition and a fixed set of properties. They cannot be physically separated into simpler substances.
• Homogeneous Matter: Matter that exhibits the same properties throughout the entire sample.   - Solutions: Homogeneous mixtures containing two or more pure substances. For example, a mixture of one spoon of sugar in a glass of water is a solution; samples taken from any part of the glass will have the identical taste and properties.
• Heterogeneous Matter: Matter with properties that vary throughout the sample. The properties of a specific sample depend on the location from which it was taken. A tomato is a classic example, as its skin, juice, seeds, and pulp possess different properties.
• Elements: Pure substances composed of homoatomic molecules or individual atoms of the same kind. Examples include oxygen gas (homoatomic molecules) and copper metal (individual copper atoms).
• Compounds: Pure substances composed of heteroatomic molecules or individual atoms of two or more different types. For example, pure water ($H_2O$) consists of heteroatomic molecules, and table salt ($NaCl$) consists of sodium ions and chlorine ions.

The Metric System and Units of Measurement

• Metric System: A decimal system where larger and smaller units are related by factors of $10$.
• Basic Unit: A specific unit from which other units for the same quantity are derived via multiplication or division.
• Derived Unit: A unit obtained by the multiplication or division of one or more basic units.
• Metric and English Units of Length:   - Metric Base: $1\,meter$   - English Base: $1\,yard$   - Metric Larger: $1\,kilometer = 1000\,meters$   - English Larger: $1\,mile = 1760\,yards$   - Metric Smaller: $10\,decimeters = 1\,meter$; $100\,centimeters = 1\,meter$; $1000\,millimeters = 1\,meter$   - English Smaller: $3\,feet = 1\,yard$; $36\,inches = 1\,yard$

Common Metric Prefixes

• mega- (M): $1,000,000 \times \text{basic unit}$ or $10^{6} \times \text{basic unit}$
• kilo- (k): $1000 \times \text{basic unit}$ or $10^{3} \times \text{basic unit}$
• deci- (d): $1/10 \times \text{basic unit}$ or $10^{-1} \times \text{basic unit}$
• centi- (c): $1/100 \times \text{basic unit}$ or $10^{-2} \times \text{basic unit}$
• milli- (m): $1/1000 \times \text{basic unit}$ or $10^{-3} \times \text{basic unit}$
• micro- (\mu): $1/1,000,000 \times \text{basic unit}$ or $10^{-6} \times \text{basic unit}$
• nano- (n): $1/1,000,000,000 \times \text{basic unit}$ or $10^{-9} \times \text{basic unit}$
• pico- (p): $1/1,000,000,000,000 \times \text{basic unit}$ or $10^{-12} \times \text{basic unit}$

Temperature Scales and Conversions

• Scales: Fahrenheit ($^{\circ}F$), Celsius ($^{\circ}C$), and Kelvin ($K$). Celsius and Kelvin are the primary scales used in scientific work.
• Formulas and Practice:   - To Fahrenheit: $^{\circ}F = \frac{9}{5}(^{\circ}C) + 32$   - To Kelvin: $K = {^{\circ}C} + 273$   - Example: Convert $22^{\circ}C$ to Fahrenheit: $\frac{9}{5}(22) + 32 = 71.6^{\circ}F \Rightarrow 72^{\circ}F$   - Example: Convert $22^{\circ}C$ to Kelvin: $22 + 273 = 295\,K$   - Example: Convert $54^{\circ}C$ to Fahrenheit: $\frac{9}{5}(54) + 32 = 129.2^{\circ}F \Rightarrow 129^{\circ}F$   - Example: Convert $54^{\circ}C$ to Kelvin: $54 + 273 = 327\,K$

Commonly Used Metric Units and Relationships

• Volume:   - $1\,dm^3 = 1\,L$   - $1\,cm^3 = 1\,mL = 1\,cc$   - $1000\,mL = 1\,dm^3$
• Mass:   - $1000\,g = 1\,kg$   - $1,000,000\,mg = 1\,kg$
• Energy:   - calorie (cal): Basic unit.   - kilocalorie (kcal): $1\,kcal = 4184\,J$   - joule (J): $1\,cal = 4.184\,J$
• Time: second ($s$) is the basic unit.
• English Equivalents:   - $1\,m = 1.094\,yd$   - $1\,cm = 0.394\,in.$   - $1\,kg = 2.20\,lb$   - $1\,L = 1.057\,qt$   - BTU (British Thermal Unit): The amount of heat required to raise the temperature of $1\,pound$ of water by $1^{\circ}F$.

Scientific Notation

• Representation: Numbers are expressed as $M \times 10^{n}$.
• Nonexponential term (M): A number between $1$ and $10$ (exclusive of $10$), written with a decimal.
• Exponential term (n): Ten raised to a whole number exponent.   - If $n$ is positive, the original decimal was to the right of the standard position (the position after the first non-zero digit).   - If $n$ is negative, the original decimal was to the left of the standard position.

Significant Figures

• Definition: Numbers in a measurement representing certainty, plus one final estimated digit.
• Guidelines for Zeros:   - Leading Zeros: Zeros not preceded by non-zero numbers are NOT significant.   - Buried/Confined Zeros: Zeros between non-zero numbers are ALWAYS significant.   - Trailing Zeros: Zeros at the end of a number are significant ONLY if a decimal point is present.
• Calculations:   - Multiplication/Division: The result must have the same number of significant figures as the quantity with the fewest significant figures.   - Addition/Subtraction: The result must have the same number of places to the right of the decimal ($prd$) as the quantity with the fewest decimal places.
• Rounding Rules:   - If the first digit to be dropped is $\ge 5$, increase the last significant figure by $1$.   - If the first digit to be dropped is $< 5$, leave the last significant figure unchanged.   - Example: Round $10.825$ to $1\,prd \Rightarrow 10.8$.   - Example: Round $-0.175$ to $1\,prd \Rightarrow -0.2$.
• Exact Numbers: These have no uncertainty and do not limit significant figures. They include defined relationships ($100\,cm = 1\,m$), counting numbers ($12\,eggs = 1\,dozen$), and simple fractions ($\frac{5}{9}$).

The Factor-Unit Method

• Step 1: Write down the known quantity including its numerical value and units.
• Step 2: Set the known quantity equal to the units of the unknown quantity, leaving space for factors.
• Step 3: Multiply by conversion factors so that the units of the factor cancel the known units and produce the target units.
• Step 4: Perform the arithmetic.
• Example: Convert $1834\,cm$ to meters.   - Relationship: $100\,cm = 1\,m$   - Calculation: $1834\,cm \times \frac{1\,m}{100\,cm} = 18.34\,m$

Percentage and Density

• Percentage: Represents the number of specific items in a group of $100$. Formula: $\% = (\frac{\text{part}}{\text{total}}) \times 100$
• Density: Defined as the mass of a sample divided by the volume of that sample. $d = \frac{m}{V}$.
• Density Example:   - Mass of empty beaker: $31.447\,g$   - Mass of beaker + $20.00\,mL$ liquid: $55.891\,g$   - Mass of liquid: $55.891\,g - 31.447\,g = 24.444\,g$   - Density: $\frac{24.444\,g}{20.00\,mL} = 1.222\,g/mL$

The Periodic Law and Table

• Periodic Law: Elements arranged by increasing atomic number exhibit similar chemical properties at regular (periodic) intervals.
• Groups (Families): Vertical columns containing elements with similar chemical properties.   - U.S. System: Uses Roman numerals and letters ($A$ or $B$).   - IUPAC System: Uses numbers $1$ to $18$.
• Periods: Horizontal rows numbered from top to bottom. There are $7$ periods in total. Elements $58-71$ and $90-103$ are placed below the main table to maintain correct period alignment.
• Examples of Location:   - Calcium ($Ca$, $20$): Group $IIA(2)$, Period $4$.   - Silver ($Ag$, $47$): Group $IB(11)$, Period $5$.   - Sulfur ($S$, $16$): Group $VIA(16)$, Period $3$.

Atomic Structure and Electron Shells

• Valence Shell: The outermost, highest-energy shell containing electrons. Atoms with the same number of valence electrons share similar chemical properties.
• Electron Occupancy Patterns:   - Beryllium, Magnesium, and Calcium all have $2$ electrons in their valence shell. Strontium ($Sr$) is predicted to also have $2$ valence electrons because it shares similar properties and belongs to the same group.   - Noble Gases (Group $18$) typically fill their shells (e.g., Neon has $8$ valence electrons).

Trends in the Periodic Table

• Metals: Found in the left two-thirds of the table. Properties include high thermal/electrical conductivity, high ductility, malleability, and metallic luster.
• Nonmetals: Found in the right one-third. Properties are opposite to metals; they are brittle/powdery solids or gases.
• Metalloids: Elements in the diagonal separation zone with properties between those of metals and nonmetals.
• Metallic Property Trend: Elements become less metallic from left to right across a period and more metallic from top to bottom down a group.
• Atomic Size Trend: Decreases from left to right across a period; increases from top to bottom down a group.
• First Ionization Energy: The energy required to remove an electron from a gaseous atom. It generally increases from left to right across a period and decreases from top to bottom down a group.
• Chemical Reactivity (Group IA): Reactivity with substances like ethyl alcohol increases as the atomic number increases (top to bottom). Bubbles form more vigorously moving from Lithium to Sodium to Potassium.

Lewis Structures and the Octet Rule

• Lewis Structures: A representation where the elemental symbol represents the nucleus and inner electrons, and dots represent valence-shell electrons.   - Number of dots for representative elements matches the Roman numeral group number (e.g., $Ca$ in $IIA$ has $2$ dots; $P$ in $VA$ has $5$ dots).
• Octet Rule: Atoms gain or lose electrons to achieve an electron arrangement identical to a noble gas, usually having $8$ electrons in the valence shell.
• Simple Ions: Atoms that acquire a net charge.   - Magnesium ($Mg$): Loses $2$ electrons to become $Mg^{2+}$.   - Oxygen ($O$): Gains $2$ electrons to become $O^{2-}$.   - Bromine ($Br$): Gains $1$ electron to become $Br^{-}$.
• Predicting Charges:   - Metals: Positive charge = Group Number.   - Nonmetals: Negative charge = $8 - \text{Group Number}$.

Ionic Compounds

• Ionic Bond: The attractive force between positive and negative ions. Formed when a metal loses electrons to a nonmetal.
• Binary Ionic Compounds: Formed from just two elements.   - The metal symbol is always written first.   - Formulas represent the minimum ratio for electrical neutrality.   - Examples:     - Sodium ($Na^+$) and Fluorine ($F^-$) $\rightarrow NaF$     - Sodium ($Na^+$) and Sulfur ($S^{2-}$) $\rightarrow Na_2S$     - Aluminum ($Al^{3+}$) and Oxygen ($O^{2-}$) $\rightarrow Al_2O_3$
• Naming Binary Ionic Compounds: Name = Metal + Nonmetal Stem + $-ide$.   - Variable Charges: Use Roman numerals for transition metals (e.g., $FeCl_2$ is iron (II) chloride; $FeCl_3$ is iron (III) chloride).   - Additional Examples: $K_2O$ (Potassium oxide), $Mg_3N_2$ (Magnesium nitride), $BeS$ (Beryllium sulfide), $AlBr_3$ (Aluminum bromide).

Formula Weights and Moles

• Formula Weight: The sum of atomic weights of the atoms in an ionic compound formula unit.
• Mole: One mole of an ionic compound contains Avogadro's number ($6.022 \times 10^{23}$) of formula units.

Polyatomic Ions and Polar Bonds

• Polyatomic Ions: Covalently bonded groups of atoms carrying a net charge.   - Common Ions:     - $NH_4^+$: ammonium     - $C_2H_3O_2^-$: acetate     - $CO_3^{2-}$: carbonate     - $OH^-$: hydroxide     - $NO_3^-$: nitrate     - $PO_4^{3-}$: phosphate     - $SO_4^{2-}$: sulfate     - $ClO_3^-$: chlorate     - $CN^-$: cyanide
• Naming with Polyatomic Ions: Metal Name + Polyatomic Anion Name.   - Examples: $KClO_3$ (Potassium chlorate), $Ca(ClO_3)_2$ (Calcium chlorate), $Ca_3(PO_4)_2$ (Calcium phosphate).
• Electronegativity: The tendency of an atom to attract shared electrons in a covalent bond.
• Bond Types:   - Nonpolar Covalent: Electrons shared equally (e.g., $H_2$, $CO_2$ due to symmetry).   - Polar Covalent: Electrons shared unequally (e.g., $HCl$, $H_2O$, $N_2O$).

Binary Covalent Compounds

• Covalent Bond: Attraction between two atoms sharing a pair of electrons to fulfill the octet rule.
• Naming Rules:   1. Name the less electronegative element first.   2. Use the stem of the second element with the suffix $-ide$.   3. Use Greek prefixes to indicate the number of atoms (mono-, di-, tri-, tetra-, penta-, hexa-, etc.).   4. The prefix mono- is never used for the first element.
• Examples:   - $SO_2$: Sulfur dioxide   - $XeF_6$: Xenon hexafluoride   - $H_2O$: Dihydrogen monoxide (water)