Biochemistry for Nurses BSN Post RN Introduction to Chemistry in Nursing
Chemistry is a fundamental science that plays a crucial role in nursing and healthcare. Understanding chemistry is essential for nurses as it forms the basis of many processes within the human body, including drug interactions, the function of enzymes, and the transport of gases in the bloodstream. This understanding helps nurses make informed decisions regarding patient care, medication administration, and management of health conditions. The aim of this section is to provide a foundational understanding of basic chemistry concepts before exploring the field of biochemistry, which focuses on the chemical processes within living organisms.
What is Chemistry?
Chemistry is the branch of science that deals with the composition, structure, properties, and reactions of matter. It is often referred to as the “central science” because it connects other natural sciences, such as physics, biology, and environmental science. In the context of nursing, chemistry helps explain how different substances interact within the body and how these interactions affect physiological processes. A solid understanding of chemistry is fundamental for grasping more advanced concepts in pharmacology, pathophysiology, and biochemistry.
Introduction to Biochemistry
Biochemistry is a branch of chemistry that studies the structure, composition, and chemical reactions of substances in living organisms. It encompasses all the chemical processes that take place in living beings, such as the synthesis and metabolism of biomolecules like carbohydrates, proteins, and fats. Biochemistry emerged as a separate discipline when scientists began to investigate how living organisms obtain energy from food or how fundamental biological changes occur during a disease. Biochemistry has numerous applications in fields such as medicine, food science, and agriculture.
Matter and Its Properties
Definition of Matter: Matter is anything that has mass and occupies space. Everything around us, including our bodies, is made of matter. In chemistry, matter can exist in three physical states: solid, liquid, or gas. Understanding these states and their properties is crucial for nurses, especially when handling different types of medication and understanding how they behave under various conditions.
Substance: A substance is a form of matter that has a specific composition and distinct properties. Substances can be classified as elements or compounds. Each substance has a unique set of characteristics that differentiate it from other substances, such as color, density, melting point, and boiling point.
Mixture: A mixture is a combination of two or more substances that are not chemically combined. Mixtures can be homogeneous (uniform composition throughout) or heterogeneous (composition varies throughout). For example, air is a homogeneous mixture of gases, while a salad is a heterogeneous mixture.
Physical and Chemical Properties of Matter
Physical Properties: Physical properties are characteristics of a substance that can be observed or measured without changing its composition. Examples include color, smell, taste, melting point, boiling point, density, and solubility. For instance, when ice melts to form water, its physical state changes from solid to liquid, but its chemical composition remains the same.
Chemical Properties: Chemical properties describe a substance’s ability to undergo changes that transform it into different substances. These properties depend on the composition of the substance. For example, when water undergoes electrolysis, it decomposes into hydrogen and oxygen gases, a process that involves a chemical change.
Elements and Their Properties
Definition of an Element: An element is a pure substance consisting of only one type of atom, characterized by its atomic number, which is the number of protons in its nucleus. Elements cannot be broken down into simpler substances by ordinary chemical means. They can exist as solids, liquids, or gases. For example, sodium, copper, and zinc are solids; mercury and bromine are liquids; nitrogen, oxygen, chlorine, and hydrogen are gases.
Classification of Elements: Elements are broadly classified into metals, non-metals, and metalloids based on their properties:
- Metals: These are elements that typically have high electrical and thermal conductivity, are malleable and ductile, and have a shiny appearance. Examples include iron, copper, and gold.
- Non-metals: These elements are usually poor conductors of heat and electricity, are brittle when solid, and lack a metallic luster. Examples include carbon, sulfur, and oxygen.
- Metalloids: Metalloids possess properties intermediate between metals and non-metals. Examples include silicon and boron.
Symbols of Elements
Elements are represented by symbols, which are abbreviations derived from their English, Latin, Greek, or German names. For example, the symbol for hydrogen is “H,” derived from its English name, while the symbol for sodium is “Na,” derived from its Latin name “Natrium.” If a symbol has one letter, it is capitalized (e.g., H for Hydrogen). If it has two letters, only the first is capitalized (e.g., Ca for Calcium).
Valency of Elements
Valency is a unique property of an element that describes its combining capacity with other elements. It is determined by the number of electrons in the outermost shell of an atom. Understanding valency is essential in predicting how elements will react with one another to form compounds. For example, the valency of sodium (Na) is 1, meaning it can form a bond with one chlorine (Cl) atom to produce sodium chloride (NaCl).
Elements and Radicals: Symbols and Common Valences
Below is a table of some common elements and radicals, their symbols, and their typical valences:
| Element/Radical | Symbol | Valency | Element/Radical | Symbol | Valency |
|---|---|---|---|---|---|
| Sodium | Na | 1 | Hydrogen | H | 1 |
| Silver | Ag | 1 | Chlorine | Cl | 1 |
| Magnesium | Mg | 2 | Bromine | Br | 1 |
| Calcium | Ca | 2 | Iodine | I | 1 |
| Barium | Ba | 2 | Oxygen | O | 2 |
| Zinc | Zn | 2 | Sulphur | S | 2 |
| Copper | Cu | 1,2 | Nitrogen | N | 3 |
| Mercury | Hg | 1,2 | Phosphorus | P | 3,5 |
| Iron | Fe | 2,3 | Boron | B | 3 |
| Aluminum | Al | 3 | Arsenic | As | 3 |
| Chromium | Cr | 3 | Carbon | C | 4 |
| Ammonium | NH4+ | 1 | Carbonate | CO3^2- | 2 |
| Hydronium | H3O+ | 1 | Sulphate | SO4^2- | 2 |
| Hydroxide | OH- | 1 | Sulphite | SO3^2- | 2 |
| Cyanide | CN- | 1 | Thiosulphate | S2O3^2- | 2 |
| Bisulphate | HSO4- | 1 | Nitride | N^3- | 3 |
| Bicarbonate | HCO3- | 1 | Phosphate | PO4^3- | 3 |
Compounds and Their Properties
Definition of a Compound: A compound is a substance made up of two or more elements that are chemically combined in a fixed ratio by mass. For example, water (H₂O) is a compound formed by the combination of hydrogen and oxygen in a fixed ratio of 2:1 by mass.
Types of Compounds: Compounds can be classified into two main types: ionic and covalent.
- Ionic Compounds: These compounds are formed by the transfer of electrons from one atom to another, resulting in the formation of ions. Ionic compounds do not exist in an independent molecular form but form a three-dimensional crystal lattice. Each ion is surrounded by oppositely charged ions, creating a strong attraction that gives ionic compounds high melting and boiling points. Ionic compounds are represented by formula units, such as NaCl (sodium chloride) and KBr (potassium bromide).
- Covalent Compounds: Covalent compounds are formed by the sharing of electrons between atoms. These compounds exist in molecular form, and their formulas are known as molecular formulas. For example, water (H₂O) and methane (CH₄) are covalent compounds.
Molecular Formula
A molecular formula represents the actual number of atoms of each element in a molecule. For example, the molecular formula of water is H₂O, indicating that each water molecule contains two hydrogen atoms and one oxygen atom.
Common Compounds and Their Chemical Formulae
Below is a list of some common compounds and their chemical formulae:
| Compound | Chemical Formula |
|---|---|
| Water | H₂O |
| Sodium chloride (Common salt) | NaCl |
| Silicon dioxide (Sand) | SiO₂ |
| Sodium hydroxide (Caustic Soda) | NaOH |
| Sodium carbonate (Washing Soda) | Na₂CO₃·10H₂O |
| Calcium oxide (Quick Lime) | CaO |
| Calcium carbonate (Limestone) | CaCO₃ |
| Sugar | C₁₂H₂₂O₁₁ |
| Sulphuric acid | H₂SO₄ |
| Ammonia | NH₃ |
Mixtures: Types and Properties
Definition of a Mixture: A mixture is a combination of two or more elements or compounds that are physically combined without any fixed ratio. Mixtures can be separated into their parent components using physical methods such as distillation, filtration, evaporation, crystallization, or magnetization.
Types of Mixtures:
- Homogeneous Mixtures: These mixtures have a uniform composition throughout. Examples include air, gasoline, and ice cream.
- Heterogeneous Mixtures: In these mixtures, the composition is not uniform throughout. Examples include soil, rock, and wood.
Differences between a Compound and a Mixture:
| Property | Compound | Mixture |
|---|---|---|
| Formation | Formed by a chemical combination of elements. | Formed by the simple mixing of substances. |
| Identity of Constituents | Constituents lose their identity and form a new substance with different properties. | Mixture retains the properties of its constituents. |
| Composition | Has a fixed composition by mass. | Does not have a fixed composition. |
| Separation | Components cannot be separated by physical means. | Components can be separated by simple physical methods. |
| Representation | Represented by a chemical formula. | Does not have a chemical formula. |
| Homogeneity | Always homogeneous. | Can be homogeneous or heterogeneous. |
| Melting Point | Has a sharp, fixed melting point. | Does not have a sharp, fixed melting point. |
Atomic Number and Mass Number
Atomic Number (Z): The atomic number of an element is the number of protons present in the nucleus of its atoms. It is represented by the symbol ‘Z.’ As all atoms of an element have the same number of protons, each element has a specific atomic number that serves as its identification number. For example, hydrogen has an atomic number of Z=1, carbon has Z=6, and oxygen has Z=8.
Mass Number (A): The mass number is the sum of the number of protons and neutrons in the nucleus of an atom. It is represented by the symbol ‘A’ and is calculated as:
A=Z+nA = Z + n
where nn is the number of neutrons. For example, a hydrogen atom has one proton and no neutrons, giving it a mass number of A=1. A carbon atom has six protons and six neutrons, giving it a mass number of A=12.
Relative Atomic Mass and Atomic Mass Unit (AMU)
Relative Atomic Mass: The relative atomic mass of an element is the average mass of the atoms of that element compared to 1/12th the mass of a carbon-12 atom. It is a dimensionless quantity that provides a comparison of the mass of different atoms on a relative scale.
Atomic Mass Unit (AMU): An atomic mass unit is defined as one-twelfth the mass of a carbon-12 atom. When expressed in grams, one atomic mass unit equals:
1 amu=1.66053906660×10−24 grams1 \text{ amu} = 1.66053906660 \times 10^{-24} \text{ grams}
Writing Chemical Formulae
Chemical formulae are symbolic representations of compounds. To write a chemical formula, follow these steps:
- Write the symbols of the elements side by side, with the positive ion (cation) first and the negative ion (anion) last.
- Indicate the valency of each ion as superscripts on the right side of the element’s symbol. For example, Na+\text{Na}^+, Ca2+\text{Ca}^{2+}, Cl−\text{Cl}^-, and O2−\text{O}^{2-}.
- Use the ‘cross exchange’ method to bring the valencies of the ions to the lower right corner of the opposite ion. If the valencies are the same, they are omitted in the final formula. For example:
Na++Cl−→NaCl\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl} Ca2++Cl−→CaCl2\text{Ca}^{2+} + \text{Cl}^- \rightarrow \text{CaCl}_2
- If an ion is a combination of two or more atoms (a radical) with a net charge, write the formula by enclosing the radical in parentheses and using the cross exchange method. For example, the formula of aluminum sulfate is written as:
Al2(SO4)3\text{Al}_2(\text{SO}_4)_3
Theories and Experiments Related to the Structure of Atom
Dalton’s Atomic Theory: According to John Dalton, an atom is an indivisible, hard, dense sphere. Atoms of the same element are identical and combine in various ways to form compounds. However, later experiments revealed that atoms consist of subatomic particles.
Discovery of Subatomic Particles: In the late 1800s and early 1900s, scientists discovered new subatomic particles:
- Protons: Discovered by Eugen Goldstein in 1886, protons are positively charged particles found in the nucleus.
- Electrons: Discovered by J.J. Thomson in 1897, electrons are negatively charged particles that orbit the nucleus. Thomson proposed the “plum pudding” model, suggesting that atoms were solid structures of positive charge with negatively charged particles embedded within, like plums in a pudding.
Atoms and Ions
Atoms: An atom is the basic unit of matter that uniquely defines a chemical element. Atoms consist of a nucleus, which contains positively charged protons and neutral neutrons, surrounded by negatively charged electrons.
Ions: An ion is an atom or a group of atoms that has an electric charge. Ions with a positive charge are called cations, while ions with a negative charge are called anions. Ions are formed when an atom gains or loses electrons. For example, sodium (Na) loses one electron to become a cation (Na+\text{Na}^+), while chlorine (Cl) gains one electron to become an anion (Cl−\text{Cl}^-).
Radicals: Also known as free radicals, radicals are atoms, molecules, or ions with at least one unpaired valence electron, making them highly reactive.
Periodic Table of Elements
The periodic table is an organized chart of elements arranged by increasing atomic number. It is based on the periodic law, which states that the properties of elements are periodic functions of their atomic numbers.
Dobereiner’s Triads: Johann Wolfgang Döbereiner observed that the atomic mass of the middle element in a group of three elements, or a triad, was approximately the average of the atomic masses of the other two elements.
Newlands Octaves: In 1864, John Newlands proposed the ‘Law of Octaves,’ which stated that every eighth element had properties similar to the first.
Mendeleev’s Periodic Table: Dmitri Mendeleev arranged the known elements (63 at the time) in order of increasing atomic mass in a table called the Periodic Table. He proposed the periodic law: “The properties of the elements are periodic functions of their atomic masses.”
Modern Periodic Table: The modern periodic table is based on the arrangement of elements according to increasing atomic numbers. Elements with similar properties are grouped together in vertical columns called groups, while horizontal rows are called periods.
Significance of the Modern Periodic Table
The modern periodic table’s significance lies in its ability to demonstrate the periodicity of elements’ properties, which are determined by their electronic configuration. The table is structured in such a way that it reflects the repeating patterns of elements’ behavior, such as their reactivity, ionization energy, and electronegativity.
Salient Features of the Modern Periodic Table
- Periods: There are seven horizontal rows called periods. Each period represents the filling of a new electron shell.
- First Period: Contains only two elements, hydrogen and helium.
- Second and Third Periods: Each contains eight elements.
- Fourth and Fifth Periods: Each contains eighteen elements.
- Sixth Period: Contains thirty-two elements.
- Seventh Period: Contains twenty-three elements and is incomplete.
- Groups: There are eighteen vertical columns called groups, numbered from 1 to 18. Elements in a group have similar chemical properties due to their similar valence electron configurations.
- Blocks: Elements are classified into four blocks (s, p, d, f) based on the type of subshell being filled with electrons.
- Lanthanides and Actinides: Two series of elements, known as lanthanides and actinides, are placed separately at the bottom of the periodic table to keep it compact.
Atomic Structure: Shells and Subshells
Shells: Electrons orbit the nucleus in specific circular paths known as shells or energy levels. Each shell is associated with a specific amount of energy. The shells are labeled K, L, M, N, etc., corresponding to energy levels 1, 2, 3, 4, etc.
Subshells: Within each shell, electrons are further grouped into subshells (s, p, d, f) based on the shape of the region they occupy. Each subshell contains orbitals where electrons are most likely to be found. For example, the s subshell can hold up to 2 electrons, the p subshell up to 6, the d subshell up to 10, and the f subshell up to 14.
Table: Shells and Subshells
| Shell/Orbit | No of Electrons | Sub Shell/Orbitals | No of Electrons |
|---|---|---|---|
| K | 2 | s | 2 |
| L | 8 | s, p | 2, 6 |
| M | 18 | s, p, d | 2, 6, 10 |
| N | 32 | s, p, d, f | 2, 6, 10, 14 |
Blocks in the Periodic Table
Definition: A block in the periodic table is a set of adjacent groups of elements with similar properties. The highest energy electrons in each element of a block belong to the same atomic orbital type. The four blocks (s, p, d, f) correspond to the type of atomic orbitals their valence electrons occupy.
Periodicity of Properties in the Periodic Table
Atomic Size and Atomic Radius: Atomic size, or atomic radius, refers to half the distance between the nuclei of two bonded atoms. For example, the distance between two carbon atoms in a molecule is 154 picometers (pm); hence, the atomic radius of carbon is 77 pm.
Trends in Atomic Size:
- Across a Period: Atomic size decreases from left to right across a period due to increasing nuclear charge, which pulls the electrons closer to the nucleus.
- Down a Group: Atomic size increases from top to bottom within a group due to the addition of more electron shells.
Shielding Effect: The shielding effect occurs when inner electrons block the attraction of the nucleus for outer electrons. As the atomic number increases, more electrons are added, increasing the shielding effect.
Ionization Energy: Ionization energy is the energy required to remove the most loosely bound electron from the valence shell of an isolated gaseous atom. It increases across a period from left to right due to increasing nuclear charge, making it harder to remove electrons. It decreases down a group as the atomic size increases, reducing the force of attraction between the nucleus and the outer electrons.
Electron Affinity: Electron affinity is the energy released when an electron is added to the outermost shell of an isolated gaseous atom. It generally increases across a period and decreases down a group.
Electronegativity: Electronegativity is the ability of an atom to attract a shared pair of electrons toward itself in a molecule. It increases across a period from left to right and decreases down a group.
Concentration Units in Chemistry
Definition: Concentration is the proportion of a solute in a solution, expressed as a ratio of the amount of solute to the amount of solvent or solution.
Types of Concentration Units:
- Percentage Mass/Mass (% m/m): Grams of solute per 100 grams of solution.
- Percentage Mass/Volume (% m/v): Grams of solute per 100 cm³ of solution.
- Percentage Volume/Mass (% v/m): Volume of solute per 100 grams of solution.
- Percentage Volume/Volume (% v/v): Volume of solute per 100 cm³ of solution.
Chemical Formula and Reactions
Chemical Formula: A chemical formula is a symbolic representation of a compound’s composition. It indicates the elements present and their relative proportions. For example, the formula for aluminum sulfate is Al2(SO4)3\text{Al}_2(\text{SO}_4)_3.
Chemical Reactions: Chemical reactions involve the formation or breaking of chemical bonds between atoms. Reactants are substances that undergo a reaction, while products are substances produced by the reaction. For example, the breakdown of hydrogen peroxide (H₂O₂) into water and oxygen can be represented by the equation:
2H2O2→2H2O+O22 \text{H}_2\text{O}_2 \rightarrow 2 \text{H}_2\text{O} + \text{O}_2
Acids and Bases
Acids: Acids have a sour taste and turn blue litmus red. They are corrosive in concentrated form. For example, hydrochloric acid (HCl) dissociates in water to produce hydrogen ions (H⁺).
Bases: Bases have a bitter taste, feel slippery, and turn red litmus blue. They are non-corrosive except for concentrated forms of sodium hydroxide (NaOH) and potassium hydroxide (KOH). Bases dissociate in water to produce hydroxide ions (OH⁻).
Redox Reactions
Definition: Redox reactions involve the transfer of electrons between reactants, resulting in changes in their oxidation states. They can be broken down into oxidation (loss of electrons) and reduction (gain of electrons) processes.
Example of Redox Reaction: In the reaction between hydrogen and fluorine:
H2+F2→2HF\text{H}_2 + \text{F}_2 \rightarrow 2\text{HF}
- Oxidation half-reaction: H2→2H++2e−\text{H}_2 \rightarrow 2\text{H}^+ + 2e^-
- Reduction half-reaction: F2+2e−→2F−\text{F}_2 + 2e^- \rightarrow 2\text{F}^-
Chemical Bonding
Definition: Chemical bonds are interactions that account for the association of atoms into molecules, ions, crystals, and other stable species. Bonds form when atoms achieve a lower total energy through their interaction.
Types of Chemical Bonds:
- Ionic Bond: Formed by the attraction between oppositely charged ions. Example: Sodium chloride (NaCl).
- Covalent Bond: Formed by the sharing of electrons between atoms. Example: Water (H₂O).
- Metallic Bond: Involves the delocalized sharing of free electrons among a lattice of metal atoms.
- Van der Waals Forces: Weak intermolecular forces, such as London dispersion forces and dipole-dipole interactions.
- Hydrogen Bonding: A special type of dipole-dipole attraction between molecules, involving a hydrogen atom bonded to a highly electronegative atom (e.g., water molecules).
Conclusion
Understanding basic concepts in chemistry is essential for nurses to comprehend how different substances interact within the human body and how these interactions impact health and disease. This foundational knowledge of chemistry paves the way for a deeper exploration of biochemistry, enabling nurses to apply scientific principles effectively in clinical practice.