Chemistry Headache

• Chemistry studies the composition of substances and their reactions.
• Chemicals compose the human body, food, and medications.
• All anatomical structures are chemicals, and physiological processes rely on chemical reactions.
• Water and electrolytes constitute all body fluids.

Fundamentals of Chemistry

• Matter is anything that occupies space and has mass, and it is composed of elements existing as solids, liquids, or gases.
• Mass is the amount of matter present.
• Weight measures heaviness based on gravitational pull on mass.
• Chemistry studies the composition, properties, and interactions of matter.
• Biochemistry studies physiological processes and disease from a biological standpoint.

Elements and Atoms

• Elements are the simplest types of matter with specific chemical properties; 92 occur naturally.
• Compounds result from chemical combinations of different elements.
• Atoms are the smallest particles of an element retaining its properties.
• The body requires different elements in varying amounts.
• Bulk elements are needed in large amounts such as carbon, oxygen, hydrogen, nitrogen, sulfur, and phosphorus (C, O, H, N, S, P).
• Trace elements are needed in small amounts such as iron and iodine (Fe, I).
• Ultratrace elements are required in very minute amounts such as arsenic (As).

Atomic Structure

• Atoms consist of subatomic particles, including protons, neutrons, and electrons.
• Protons are large particles carrying a single positive charge.
• Neutrons are large particles with no electrical charge.
• Electrons are small particles carrying a single negative charge.
• Atoms have a central nucleus containing protons and neutrons, with electrons constantly moving around it.
• The number of protons equals the number of electrons, making atoms electrically neutral.

Atomic and Mass Number

• Atomic Number signifies the number of protons in an atom's nucleus, unique to each element.
• Mass Number is the sum of protons and neutrons in an atom.
• Electrons do not significantly contribute to an atom's mass.

Isotopes

• Isotopes are atoms of an element with the same atomic number but different mass numbers.
• Isotopes have varying numbers of neutrons but the same number of protons and electrons.
• Oxygen forms isotopes like O-16, O-17, and O-18, where the numbers indicate mass numbers.
• Some isotopes are radioactive, and they release energy or atomic fragments until stable; they are harnessed in medicine for detection and treatment.
• Atomic weight is the average of the mass numbers of all isotopes of an element.

Isotopes in Medicine

• Radioactive iodine-131 is used to destroy cancerous thyroid tissue because the thyroid uniquely absorbs and metabolizes iodine.
• Radioactive isotopes are used in medicine to detect coronary blood vessel disorders, evaluate kidney function, measure hormone levels, and assess bone density changes.

Molecules and Compounds

• A molecule forms when two or more atoms combine chemically.
• A compound forms when two or more atoms of different elements combine chemically.
• Molecular formulas show the types and numbers of atoms in a molecule:
  • H₂ is hydrogen
  • C6H12O6 is glucose
  • H₂0 is water

Bonding of Atoms

• Chemical bonds occur when atoms combine, involving interactions between electrons.
• Electrons occupy spaces called electron shells around the nucleus.
• For atoms up to atomic number 18:
  • The first shell holds up to 2 electrons.
  • The second shell holds up to 8 electrons.
  • The third shell holds up to 8 electrons.
• Lower energy shells fill before outer ones, stabilizing with certain electron counts in the outermost shell (2, 8, or 18).
• Outermost electrons determine if an atom will form chemical bonds.

Ionic Bonds

• An ion is an electrically charged atom that gains or loses electrons to achieve stability.
• A cation is a positively charged ion formed when an atom loses electrons.
• An anion is a negatively charged ion formed when an atom gains electrons.
• Ionic bonds are strong chemical bonds between oppositely charged ions.

Covalent Bonds

• Covalent Bonds are strong chemical bonds, which are created between atoms that share electrons.
• Two hydrogen atoms (H) can form a stable hydrogen molecule (H₂) through covalent bonds.
• Hydrogen molecules (H₂) and oxygen molecules (O₂) combine through covalent bonds to form water (H₂O) molecules.

Nonpolar and Polar Covalent Bonds

• Nonpolar covalent bonds share electrons equally.
  • They occur between atoms with similar electronegativity.
  • Atoms of the same element share electrons equally due to the same number of protons.
• Polar covalent bonds share electrons unequally.
  • They occur between atoms with different electronegativity.
  • These bonds form polar molecules with uneven charge distribution.
  • Atoms with more protons have higher electronegativity, pulling electrons closer to the nucleus.
  • Water is a polar molecule.

Structural Formulas

• Structural formulas display atom arrangement and bonding in molecules.
• A single line between atoms indicates one pair of electrons shared (single bond).
• Two lines indicate two pairs shared (double bond).

Hydrogen Bonds

• Hydrogen bonds are weak attractions, where a slightly positive hydrogen (H) end of one polar molecule is attracted to a slightly negative nitrogen (N) or oxygen (O) end of another polar molecule.
• These bonds form between adjacent water molecules.
• They are important for the structure of proteins and nucleic acids.

Chemical Reactions

• Chemical reactions occur when chemical bonds break or form between atoms, ions, or molecules.
• Reactants are initial substances in a chemical reaction: atoms, ions, or molecules.
• Products are substances that are created toward the end of the reaction.

Types of Chemical Reactions

• Synthesis Reaction: Complex chemical structures are formed from simpler ones, as in A + B → AB.
• Decomposition Reaction: Chemical bonds are broken to form simpler structures, as in AB → A + B.
• Exchange Reaction: Chemical bonds break and new bonds form, as in AB + CD → AD + CB.
• Reversible Reaction: The products can convert back to the reactants, as in A + B ↔ AB.

Electrolytes, Acids and Bases, and Salts

• Electrolytes release ions in water, creating a solution capable of conducting electricity.
  • NaCl → Na⁺ + Cl⁻ is an example.
• Acids are electrolytes that release hydrogen ions (H⁺) in water.
  • HCl → H⁺ + Cl⁻ is an example.
• Bases release ions that combine with hydrogen ions (H⁺).
  • NaOH → Na⁺ + OH⁻ is an example.
• Salts are electrolytes formed from reactions between an acid and a base.
  • HCl + NaOH → H₂O + NaCl is an example.

Acids and Bases

• Measuring H⁺ ion concentration in a solution determines acidity or alkalinity.
• Ionically bonded substances in water attract the slightly charged ends of polar water molecules.
• This process dissociates the substance into ions surrounded by water, creating an electrolyte capable of carrying electric current.

pH and H⁺ Concentration

• The pH scale measures the concentration of hydrogen ions in a water-based solution.
• A pH of 7 indicates a neutral solution where H⁺ and OH⁻ concentrations are equal, which is the pH of water.
• Acidic solutions have a pH less than 7, signifying a greater concentration of H⁺ than OH⁻.
• Basic (alkaline) solutions have a pH greater than 7, indicating a higher concentration of OH⁻ than H⁺.

pH Scale Details

• The pH scale ranges from 0 to 14.
• Each pH unit represents a tenfold difference in H⁺ ion concentration.
• Acids have a pH <7, and bases have a pH >7.
• A pH of 7 is neutral.

Conditions due to Abnormal pH

• Normal blood pH ranges from 7.35 to 7.45.
• Acidemia occurs when blood pH drops to 7.0-7.3, which causes acidosis symptoms like disorientation and fatigue, and could be caused by Alkaline intestinal contents, lung disease or diabetes.
• Alkalemia occurs when blood pH rises to 7.5-7.8, which causes alkalosis symptoms like dizziness and agitation, and can be caused by high altitude breathing, high fever, or excess antacids.
• Homeostatic mechanisms regulate pH, and buffers counteract pH changes by binding or releasing H⁺ ions.

Organic vs. Inorganic

• Organic molecules contain carbon and hydrogen.
  • Solubility depends of the compound, can ether be water or organic liquids
  • Carbohydrates, proteins, lipids, and nucleic acids are examples of organic materials.
  • Water-soluble compounds do not release atoms and are non-electrolytes
• Inorganic molecules generally do not have carbon.
  • Water, salts, and carbon dioxide are examples.
  • Typically dissolve in water, forming electrolytes.

Inorganic Substances

• Water makes up two thirds of the the organisms body
  • It acts as a medium for transport in the body, and for the transport across most metabolic reactions
  • Water balance occurs when intake equals loss
• Oxygen (O₂) is used by organelles to release energy from nutrients, powering the cell's metabolic functions.
  • Survival necessitates oxygen.
• Carbon dioxide (CO₂) is a waste product from metabolic reactions and releasing into the atmosphere as one exhales.
• Inorganic Salts:
  • Abundant in H2O, which is bodily fluid
  • They are sources of Na+, Cl-, K+, Ca², etc

Organic Substances: Carbohydrates

• Carbohydrates act as a main energy source for cellular activities
  • Supply materials to build cell structures, are soluble in water and containing C, H, and O.
  • Ratio of elements usually 1:2 (C, H.O), like (C6H12O6 = glucose) Sizes:
  • Monosarrides are single sugars such as fructose and glucose
  • Disaccharides are double sugars such as lactose and sucrose
  • Polysaccharides are complex such as cellulose ansd glycogen

Organic Substance, Lipids

• Lipids are insoluble in water and are soluble in non-polar solution
• Lipid molecules consist of triglycerides are fats and phospholipids, and steroids
• Lipids component of cell membranes, and have various other functions withhin the cells
• Abundant lipids consist of triglycerides that:
  • Are used for cellular energy
  • Carry more energy in the cell, per-gram
  • But less (O) than carbohydrates
  • 1 Glyerol mlcl with 3 fatty acids

Fatty Acids

• Have only single carbon bonds and tend to be of animal origion
• Whereas unsaturated have one or more double or triple bonds between the carbon
• Liquids are room temperature, and extracted from plants

Phospholipids

• Glycerol that consists of 2 fatty acids and 1 phosphate acid

Steroids

• Consists of 4 carbon rings, that is distrubuted in the in different functions
• A key component in synthesizing adrnal, and synthesis sex hormones,
• Cholesterol is main steroid in of the body

Organic Substances: Proteins

• Proteins serve as structural components and act as receptors, enzymes, and energy sources
• Proteins consist of amino acids as a building block
• Amino acids (-NH2), carboxyl {COOH), and a an unique R group
• Amino acids are bound to each other by bonds and amino acids

Protein Structure

• Primary : Amino acid sequence
• Secondary: Twistesr structure that is formed by h-bonds
• Teritary" Unique in shapes and dimensions, helps create and create a function
• Quaterniary: Formed by more than one chain connected to a Protein

Proteins: Denaturation

• Which is cause by the temperature and PH shift in change
  • Irreversible cases can effect enzymes, and create failure to function

Organic Substances: Nucleic Acids

• Carry genetic code (DNA) or aid in protein synthesis (RNA), encoding amino acid sequences of proteins. A building block, nucleoides: consists of a sugar (S), that has phosphate group
• DNA Is double chain, as RNA is a single chain
  • RNA Interactd with DNA to aid Protein

Nucleic Acids

• Have a a double helix form
• Have deoxyribose And Composed by nucleotides

RNA

• Interacting with DNA to help with protein synthesis
• Have a structural strand-single helix
• Composed of nucleotides

Science For Tech: Imaging

• Used to visualize internal anatomy
  • CT scans involve radiation for visualizing the internal organs
  • PET scans use unusual radioactive isotopes to detect bio activity