Chemistry I (Basic and Inorganic) PDF

# Chemistry I (Basic and Inorganic) As found in: Lecture Notes for Human Anatomy & Physiology – Bio 32 (Newton 2020) Figure Credit: Chapter 2 Marieb & Hoehn, 2020. Human Anatomy & Physiology (7th Ed.) (unless noted otherwise) ## Lecture Objectives - Define chemical element and list the four ele...

# Chemistry I (Basic and Inorganic) As found in: Lecture Notes for Human Anatomy & Physiology – Bio 32 (Newton 2020) Figure Credit: Chapter 2 Marieb & Hoehn, 2020. Human Anatomy & Physiology (7th Ed.) (unless noted otherwise) ## Lecture Objectives - Define chemical element and list the four elements that form the bulk of body matter. - List the subatomic particles and describe their relative masses, charges, and positions in the atom. - Draw the atomic structure of any given atom. - Define atomic number, atomic mass, isotope, and radioisotope. - Explain the role of electrons in chemical bonding in relation to the octet rule. - Name and describe four types of chemical bonds: ionic, polar covalent, nonpolar, covalent, and hydrogen bonds. - Distinguish between organic and inorganic compounds. - Differentiate between a salt, an acid, and a base. - Explain the importance of water to body homeostasis and provide several examples of the role water plays in the human body. - Explain the concept of pH and state the pH of blood. ## Key Terms | Term | Description | |---|---| | Matter | Anything that occupies space and has mass. | | Atom | The smallest unit of matter that still retains the properties of an element. | | Element | Substances that cannot be broken down into substances with different properties. Composed of one type of atom.| ## What is matter? - Anything that occupies space and has mass. - Physical “stuff” (living and non-living) of the universe - Matter is found on the Earth in three physical states: - Solid - Liquid - Gas - Mass vs. Weight ## Composition of Matter - Twenty-five elements are essential to life - Four elements make up about 96% of the human body: - Hydrogen - Oxygen - Nitrogen - Carbon The image depicts a human figure with the following percentages of each element: - Oxygen (O): 65.0% - Carbon (C): 18.5% - Hydrogen (H): 9.5% - Nitrogen (N): 3.3% The remaining elements make up less than 0.01% of the human body and are labeled as “trace elements” including: - Boron (B) - Chromium (Cr) - Manganese (Mn) - Molybdenum (Mo) - Cobalt (Co) - Copper (Cu) - Fluorine (F) - Iodine (I) - Iron (Fe) - Selenium (Se) - Silicon (Si) - Tin (Sn) - Vanadium (V) - Zinc (Zn) ## Trace Element Deficiency The image depicts a woman with a swollen neck, highlighting the importance of trace elements in human health. ## Periodic Table of the Elements The image depicts a standard periodic table of the elements, with information about each element including atomic number, symbol and approximate atomic mass. ## Atoms - Building blocks of elements - Atomic Structure - Nucleus - Protons (p+) - Neutrons (n°) - Outside of nucleus - Electrons (e-) The image depicts a simple model of an atom with the following elements labeled: - Nucleus - 2 Electron cloud - 2 Protons - 2 Neutrons - 2 Electrons ## Identifying Elements - Elements differ in the number of subatomic particles in their atoms. - Each element has a different chemical (atomic) symbol - The number of protons, the atomic number, determines which element it is. - An atom’s mass number (atomic mass) is the sum of the number of protons and neutrons. The image depicts a simple diagram of a Carbon atom, indicating the following information: - Atomic Number: 6 - Mass Number: 12.011 - Chemical symbol: C ## Isotopes - Same number of protons (+) and electrons (-) - Vary in number of neutrons. - Ex: 12C, 13C, 14C - Radioisotopes - Heavy isotope. - Tends to be unstable. - Decomposes more than stable isotope. - Radioactivity - Process of spontaneous atomic decay - As some _isotopes_ adjust to a more stable form, they will emit a measurable energy. This energy release is called “radiation”. - We can make use of radioactive isotopes in medicine (low level radiation) The image depicts two applications of radioactive isotopes: - X-rays: The image shows a visual representation of an X-ray scan, revealing the shape of the thyroid gland. - CT scan: The image shows a computer tomography scan, highlighting its use in medical imaging. ## Electrons and Chemical Bonding - Electrons determine how an atom behaves when it encounters other atoms. - Electrons orbit the nucleus of an atom in specific electron shells - First electron shell (can hold 2 electrons) - Outer electron shell (can hold 8 electrons) The image shows a diagram of electron shells of four different elements: - Hydrogen (H), Atomic Number: 1 - Carbon (C), Atomic Number: 6 - Nitrogen (N), Atomic Number: 7 - Oxygen (O), Atomic Number: 8 - Bonding involves interactions between electrons in the outer shell (valence shell). - Full valence shells do not form bonds (stable) - Inert elements - Have complete valence shells and are stable. - Ex. He and Ne The image shows diagrams of the electron shells of the two inert elements: - Helium (He), Atomic Number: 2 - Neon (Ne), Atomic Number: 10 - Reactive elements - Valence shells are not full = unstable - Tend to gain, lose or share electrons (ionic bonding) - Allows for bond formation, which produces stable valence - Ex. H, O, N, C The image shows diagrams of electron shells of the four reactive elements: - Hydrogen (H), Atomic Number: 1 - Carbon (C), Atomic Number: 6 - Oxygen (O), Atomic Number: 8 - Sodium (Na), Atomic Number: 11 The image shows detailed diagrams of the electron shells of the five elements: - Hydrogen (H), 1p+, 0n, 1e - Oxygen (O), 8p+, 8n, 8e - Sodium (Na), 11p+, 12n, 11e - Carbon (C), 6p+, 6n, 6e ## Key Terms | Term | Description | |---|---| | Ionic Bond | Attraction between atoms of opposite charges (transfer of electrons). | | Covalent Bond | Sharing of electrons between atoms. Considered the strongest bond. | | Hydrogen Bond | Attraction between partially charged (polar) molecules. Considered the weakest bond. | ## Chemical Bonds - Ions - atoms in which the number of electrons _does not equal_ the number of protons - Charged particles - Ex. NaCl - Will dissociate in water and form + and – ions. Image depicts the formation of sodium ions (Na+) from sodium atoms (Na) highlighting the loss of an outer electron. - The outer electron is stripped from sodium and completes the chlorine atom's outer shell - The attraction between the ions—an ionic bond—holds them together The image depicts the formation of sodium chloride (NaCl), highlighting the transfer of an electron from Sodium (Na) to Chlorine (Cl). - The properties of a _compound_ differ from those of its atoms. The image depicts the transformation of sodium (silvery metal) and chlorine (poisonous gas) into sodium chloride (table salt). ## Covalent Bonds - A covalent bond forms when two atoms share one or more pairs of outer-shell electrons. - Atoms held together by covalent bonds form a _molecule_. - Strongest bond **✓** - The number of covalent bonds an atom can form is equal to the number of additional electrons needed to fill its outer shell. - Single covalent bonds share _one_ electron - Double covalent bonds share _two_ electrons The image depicts the formation of the following molecules through covalent bonds: - Hydrogen gas (H2) - Oxygen gas (O2) - Methane gas (CH4) The image also depicts the following representations of the same molecules: - Electron configuration - Structural formula - Space-filling model - Ball-and-stick model - Polarity - Covalent bonded molecules - Nonpolar bonds = _shared_ - Electrically neutral as a molecule as electrons are evenly shared - Ex. CO2 - Polar _unevenly shared_. - Electrons are unevenly shared resulting in a positive and negative side - Ex: H2O The image depicts the structure of carbon dioxide (CO2), highlighting the even sharing of electrons in the covalent bonds. ## Hydrogen Bonds - Weak chemical bonds **✓** - Hydrogen is attracted to the negative portion of the polar molecule. - Provides attraction between molecules. The image depicts the hydrogen bonds between multiple water (H2O) molecules. - The electronegative oxygen atom of one water molecule will form a hydrogen bond with the slightly positive hydrogen atom of another water molecule The image shows a humorous depiction of a water molecule, emphasizing that the oxygen atom is electronegative, attracting the slightly positive hydrogen atoms of another water molecule. - The hydrogen bonds provide attraction between water molecules. The image shows a water strider standing on the surface of water, demonstrating how the surface tension produced by hydrogen bonding allows the water to support its weight. ## Bond Strength - Covalent > Ionic > Hydrogen ## Molecules vs. Compounds - Molecule = particles consisting of two or more atoms held together by chemical bonds. - Usually same element, but not always - When talking about a molecule of the same element, connected via covalent bond - Examples: O2, CH4, N2, H2, H2O - Oxygen and nitrogen - Compound = substance composed of two or more _different_ elements held together by chemical bonds. - Examples: CH4, H2O, NaCl, C6H12O6,. - H2O, H2, CH4, NaCl - All compounds are molecules, not all molecules are compounds. _Element_ - Note that because electrolytes are electrically attracted and not chemically combined, we do not use the term “molecule" to describe NaCl etc. ## Solutions - Solutions: two or more components physically intermixed (not chemically bound). - Ex: saline solutions (table salt (NaCl) and water), blood plasma, interstitial fluid, urine, etc. - Solvent: dissolving medium - Present in greatest amount - Ex: water (body’s best example) - Solute: the dissolved substance - Ex: NaCl, glucose, O2, CO2, Ca2+, etc. - Present in smaller amounts. The image depicts the process of dissolution of salt in water, highlighting the distinction between solvent and solute. ## Concentrations - The amount (concentration) of a solute in the total solution is usually measured as one of the following: - Percent of the solute in the total solution (parts per 100 parts) - Milligrams per deciliter (mg/dl) - A deciliter is 100 mL - Molarity (moles per liter), indicated by M ## Key Terms | Term | Description | |---|---| | Reactants | Reacting substances. | | Products | End product (result). | | Metabolism | Sum of all chemical reactions. | The image illustrates the differences between reactants and products in a chemical reaction. ## Chemical Reactions - 4H + C → CH4 - Reactants - Product The image depicts a general chemical reaction equation, highlighting the direction of change from reactants to products with the use of an arrow. ## Types of Chemical Reactions - Synthesis reactions: Smaller particles are bonded together to form larger, more complex molecules. - Decomposition reactions: Bonds are broken in larger molecules, resulting in smaller, less complex molecules. - Exchange reactions: Bonds are both made and broken (also called displacement reactions). The image depicts the differences between three types of chemical reactions: - Synthesis reactions: Show the bonding of amino acids to form a protein molecule. - Decomposition reactions: Show the breakdown of glycogen molecules into glucose molecules. - Exchange reactions: Show the transfer of a phosphate group by ATP to glucose, resulting in the formation of glucose phosphate and ADP. ## Inorganic Compounds - Lack Carbon - Exceptions: CO2, HCO3- - Tend to be simpler than organic compounds - Ex: H2O, NaCl (inorganic) vs. C 6H12O6 (organic) The image depicts a simplified distinction between inorganic and organic compounds ## Important Inorganic Compounds in Living Matter ### Water - H2O - Most abundant inorganic compound (60-80% of our body weight) - Vital properties - High heat capacity - High heat of vaporization - Polar solvent properties - Often called the “universal solvent" - Biological molecules do not react unless they are in solution, so this is crucial to sustain life - Will also form hydration layers to shield charges - Serves as a transport medium - Chemical reactivity - Hydrolysis reactions: add water molecule - Dehydration synthesis: remove water molecules - Cushioning - Stabilizes structure of all macromolecules - Hydrophobic interactions ### Oxygen - O2 - 20% of the air you breathe is oxygen. - O2 is essential for cells to extract energy from other compounds. - Without O2, many cells die quickly as they run out of internal energy compounds (ATP). ### Carbon Dioxide - CO 2 - As energy is extracted from molecules with long chains of carbon atoms (_bonds_ are broken, and carbon atoms must be removed from the body. CO 2 is formed. ### Salts - Ionic compounds that contain cations other than H+ and anions other than OH- - Ex: NaCl, CaCO3, KCI, etc. - Easily dissociate into ions in presence of water - Vital to many body functions - Nerve cell communication, muscle contraction, etc. - If ionic balance in our body is not maintained (a function of the kidneys), the physiological activities listed above, and thousands of others will become disrupted and stop. Virtually nothing in the body will then work and death will quickly ensue. - Lithium Chloride - LiCl + H2O → Li+ + Cl- The image shows a simplified visual representation of the dissociation of NaCl into ions. ## Acids and Bases - Acids - Proton donors - (H+) = proton - Electrolytes are called acids if they yield H+ in water - Ex of partial disassociation: (weak acid) H2CO3 - (note: the body uses the decomposition part of this reversible reaction to form CO 2 for removal from the body) - H2CO3 → H⁺ + HCO3⁻ - Ex of complete disassociation: (strong acid) HCI - (note: the stomach uses this disassociation to create a very acidic stomach environment) - HCl → H⁺ + Cl⁻ - Bases - Proton acceptors - They attract and combine with H+ in water - examples: - NH3 + H+ → NH4+ - H+ + HCO3⁻ → H2CO3 - The kidney uses the synthesis reaction above to get rid of excess acid (H+) ## pH - Measures H+ (proton) concentration - Scale runs from 0-14 - pH = 7 is considered neutral - pH < 7 = acid - pH > 7 = base - Concept of physiological pH - Normal pH range for human blood is 7.35-7.45 - Therefore, <7.35 = acidosis and >7.45 = alkalosis in humans - Buffers - chemicals that can regulate pH change - ex: carbonic acid-bicarbonate system - CO2 + H2O → H2CO3 → H+ + HCO 3 ## Solutions - Are these solutions acidic, basic, or neutral? Explain why. The image shows three glasses with different number of H+ and OH- ions, and labels them as: - Basic - Neutral - Acidic - The pH scale and pH values of representative substances. The image represents a pH scale with values ranging from 0 to 14 and a color bar highlighting the acidic and basic ranges along with examples of common substances and their pH values.