Chemistry Chapter: Bonds and Solutions

Molecules of Life 1

• Textbook chapters cover topics 2-4
• Focuses on the chemistry of life
• Includes valence, covalent bonds, noncovalent bonds (hydrogen bonds), basic organic chemistry, molecular shape, and isomers.
• Review of atoms, ions, protons, electrons, neutrons, elements, and compounds is assumed but will be reviewed.

Chemical Composition of a Human Being

• Human beings are primarily composed of oxygen (65%), followed by carbon (18.5%), hydrogen (9.5%), and nitrogen (3.3%) by weight.
• The relative abundance of elements in the Earth's crust is different from that in living organisms.
• Exam questions will not focus on the percentage of elements in the human body or Earth's crust.

Development from a Single Cell

• All humans develop from a single cell.
• The human body is comprised of trillions of cells with varied functions
• PoB1 will detail how simple building blocks make cells, how these cells communicate, and how they cooperate to form organs and organisms.

Chemistry Knowledge Assumed

• Understanding of atoms, ions, protons, electrons, and neutrons is presumed.
• Knowledge of elements and compounds is assumed.

Determining Compound Properties

• A compound's properties depend on the arrangement of atoms (structure) and how atoms are bonded.
• The number of protons determines an atom's identity.
• Electron distribution plays a key role in forming bonds.

Elements and Compounds

• Matter is composed of elements.
• An element cannot be broken down into simpler substances by chemical reactions.
• A compound is a substance formed from two or more elements in a fixed ratio.
• Compound properties differ from its constituent elements.

The Elements of Life

• 92 natural elements exist, but only 20-25 are essential for life in humans.
• Oxygen, hydrogen, carbon, and nitrogen make up about 96.3% of living matter.
• Calcium, phosphorus, potassium, and sulfur make up the remainder of the crucial 4%.
• Trace elements are also necessary but in minute quantities.

Properties of Elements

• Each element is composed of unique atoms.
• An atom is the smallest unit of matter retaining element properties.
• Atoms are composed of subatomic particles—neutrons (neutral), protons (positive), and electrons (negative).
• The nucleus contains protons and neutrons; electrons orbit the nucleus.

Electron Distribution and Chemical Properties

• Electrons reside in electron shells with varying energy levels.
• An atom's chemical behavior depends on electron shell distribution.
• The periodic table organizes elements by electron distribution.

Properties of Elements

• Mass number (A): Represents the sum of protons and neutrons.
• Atomic number (Z): Is equal to the number of protons in an atom’s nucleus.

Periodic Table

• Displays known elements ordered by increasing atomic number.
• Elements in each row exhibit electron addition patterns.

Valence

• Valence electrons are outermost shell electrons.
• Atomic chemical behavior is predominantly determined by valence electron numbers.
• Atoms with full valence shells are chemically inert.

Goal of Valence Shells

• Atoms interact to complete their valence shells.
• First shell holds a maximum of 2 electrons.
• Second and third shells hold a maximum of 8 and 18 electrons, respectively.

Poll Question

• Atoms containing 8 electrons in their outer shell tend to be stable.

Chemical Reactions

• Chemical reactions involve the making and breaking of chemical bonds.
• Reactants are initial molecules.
• Products are resultant molecules.

Chemical Bonding

• Atoms form bonds with incomplete valence shells to attain stability.
• Covalent bonding involves electron sharing.
• Ionic bonding involves electron transfer.
• Hydrogen bonding results from attractions between polar molecules.

Types of Chemical Bonds

• Covalent bonds, Ionic bonds, Hydrogen bonds, Van der Waals attractions

Covalent Bonds

• Covalent bonds link atoms within molecules; they can form among the same or different elements.
• Bonds must add up to an atom's valence.
• A single bond shares one electron pair, a double bond shares two, and a triple bond shares three.

Notation

• Molecular formula represents a molecule's composition.
• Condensed formula is a concise representation, often omitting hydrogen details.
• Structural formulas show atom connections and bonding locations.
• Line diagrams or skeletal formulas offer a simplified illustration method, often used for complex molecules.

Electronegativity

• Atoms in a molecule differently attract electrons (electronegativity).
• Electronegativity is the tendency of an atom to pull shared electrons closer..
• Higher electronegativity means more electron pulling power

Types of Covalent Bonds

• Nonpolar covalent bonds feature equal electron sharing.
• Polar covalent bonds feature unequal electron sharing. The more electronegative atom results in a partial negative (δ⁻) charge; the less electronegative atom has a partial positive (δ⁺) charge.

Polar Covalent Bonds

• Electrons in polar covalent bonds predominantly reside near the more electronegative atom.

Nonpolar Covalent Bonds

• Nonpolar covalent bonds are exemplified by atoms such as Hydrogen (H₂), Chlorine (Cl₂), Oxygen (O₂), and Nitrogen (N₂).

Ionic Bonds

• Ionic bonds involve electron transfer between atoms.
• Positively charged ions are cations; negatively charged ions are anions.
• Attraction between anions and cations forms ionic bonds.

Ionic Bonds

• Ionic compounds are formed when ionic bonds form.
• Salts such as sodium chloride (table salt) are often found as crystals.
• Salting is a common example of ionic bonding, occurring naturally as crystal lattices.
• Ionic compounds are stable in dry conditions but readily dissociate in water.

Hydrogen Bonds

• Hydrogen bonds form when a hydrogen atom already covalently linked to one electronegative atom is attracted to another electronegative partner.
• Common in living cells, mostly with oxygen or nitrogen.
• Water's polarity enables hydrogen bonding among its molecules.

Water as a Solvent

• Water dissolves many substances because of its polarity.
• A solution is a homogeneous mixture of substances wherein a solvent dissolves solutes to form a solution, making water the most important solvent for life.
• When an ionic compound dissolves in water, each ion becomes surrounded by water molecules, forming a hydration shell.
• Water can dissolve compounds formed from non-ionic polar molecules as well. Polar molecules can have regions containing ionic and polar regions that easily dissolve in water

Water (Hydrophilic and Hydrophobic Substances)

• Hydrophilic substances are water-attracting; hydrophobic substances repel water.

Nonpolar Molecules and Water

• Oil and similar nonpolar molecules are hydrophobic and cluster in water, minimizing the water-oil area of contact.

Acidic and Basic Conditions

• Water molecules dissociate and reform in dynamic equilibrium, with equivalent H⁺ and OH⁻ concentrations.
• Acids increase H⁺ concentration; bases reduce H⁺ concentration.

pH Scale

• pH measures H⁺ concentration in a solution.
• Acidic solutions have a pH less than 7; neutral solutions have a pH of 7; basic solutions have a pH greater than 7.

Van der Waals Interactions

• Unbalanced electron distribution creates temporary partial charges, leading to weak attractions among molecules which are close together

Molecular Shape and Function

• A molecule's size and shape influence its function.
• Atomic orbital positions determine molecular shapes.
• Hybridization of atomic orbitals shapes molecules.

Carbon as a Foundation

• Carbon forms the essential backbone of biological molecules.
• Carbon's ability to form four bonds is fundamental for its roles in molecules.

Organic Chemistry Properties

• Carbon has a valence of 4.
• Carbon bonds to four other atoms, creating complex structures.
• Examples of carbon compounds include methane (CH₄), ethane (C₂H₆), and ethene (ethylene, C₂H₄).

Drawing Organic Structures

• Chemists often implicity represent carbon and hydrogen connections in complex organic molecules, using corners for carbon and typically omitting hydrogens except in simple examples

Isomers

• Isomers are molecules with the same molecular formula but different structures and properties.
• Structural isomers differ in bond connectivity.
• Geometric isomers, also known as cis-trans isomers, differ in spatial arrangements around double bonds.

Structural Isomers

• Molecules share the same molecular formula but have different atomic connections

Geometric Isomers

• Geometric isomers have identical atom connections but differ in spatial arrangements around double bonds, which cannot rotate.

Geometric Isomers

• Multiple C=C double bonds can lead to complex structural variations in large molecules.

Tetrahedral Structure

• Carbon atoms with four single bonds have a tetrahedral shape.

Chiral Carbons and Stereoisomers

• Chiral carbon atoms are asymmetric.
• Chiral carbon atoms are bonded to four different functional groups.
• Stereoisomers (mirror image molecules) cannot be superimposed and have different properties.

Why Care About Stereoisomers?

• Many biological molecules contain chiral carbons.
• Biological systems are highly sensitive to molecular shape differences (as with stereoisomers).
• For instance, different isomers of the same molecule have vastly different properties, e.g., carvone (smells like spearmint or caraway) or ibuprofen (painkiller).

Test your knowledge on covalent and ionic bonds, their properties, and the behavior of solutions in chemistry. This quiz explores essential concepts like cell communication, isomers, and the pH scale, crucial for understanding chemical interactions. Dive into how these principles apply to real-world scenarios, especially focusing on water's unique behavior as a solvent.