Lec 1: Chem Basics - Flashcards

The Scientific Method

• Definition of the Scientific Method: The scientific method is the systematically defined process used by scientists and clinical professionals to gather and interpret information.

• Scientific Laws:

  • Scientific laws are statements that describe natural phenomena that are consistently and reproducibly observed.

  • Example: The Law of Gravity originated when Sir Isaac Newton consistently and reproducibly observed apples falling from a tree.

• Explaining Observations:

  • Explaining observations is a central component of the scientific method.

  • The process begins with the construction of a hypothesis, defined as a tentative explanation or educated guess based on presently known facts.

  • Clinical Example: If a patient complains of stomach pains, the clinician will ask a few questions and make preliminary observations before formulating a hypothesis (diagnosis) regarding the nature of the health problem.

• Testing Hypotheses:

  • The most critical phase of the scientific method occurs after a hypothesis has been constructed: it must be thoroughly tested by conducting careful experiments.

  • Experiments must be designed specifically so that all observations made are directly related to the question at hand.

  • Example: If a patient exhibits stomach pains, ordering an X-ray of his big toe will not provide relevant observations to identify the cause of the illness.

• Scientific Theories:

  • Once a hypothesis has survived extensive and repeated testing, it may be elevated to a theory—an experimentally tested explanation of an observed behavior.

  • For a theory to remain valid, it must meet three fundamental criteria:

    1. It must be consistent with all existing experimental evidence.

    2. It must accurately predict the results of future experiments.

    3. It must successfully explain future observations.

Classification of Matter

• Definition of Matter: Matter is defined as anything that takes up space and has mass (weighs something).

• Physical Properties of Matter:

  • Physical properties are characteristics that can be determined without altering the chemical composition of matter (what the substance is made of).

  • Example of Physical Properties: Sugar is physical described as white, sweet-tasting, and odorless. Measuring its melting point does not transform sugar into a different chemical substance.

• Primary Categories of Matter: Matter is classified into two main categories: pure substances and mixtures.

Pure Substances

• Definition: A pure substance is matter that is composed of only one type of substance and can be uniquely represented with a single chemical formula or symbol.

• Sub-classifications of Pure Substances:

  • Elements: The simplest type of matter, consisting of only one type of atom.

    • Atom: The smallest fundamental unit of matter that retains its unique physical and chemical characteristics.

  • Compounds: Pure substances formed from two or more different elements that are chemically joined (bonded) together.

Mixtures

• Definition: A mixture is a physical combination of two or more distinct substances. Mixtures can be physically separated into their individual components.

• Sub-classifications of Mixtures:

  • Homogeneous Mixtures: Mixtures whose chemical and physical composition is completely uniform and identical throughout.

  • Heterogeneous Mixtures: Mixtures whose composition is non-uniform and varies from one region of the sample to another.

Practice Problems: Classifying Matter

• Classification as Pure Substance or Mixture:

  1. Cake batter: Mixture.

  2. The helium gas inside a balloon: Pure substance (Element).

• Classification of Mixtures as Homogeneous or Heterogeneous:

  1. Olive oil: Homogeneous mixture (uniform composition throughout).

  2. Rocky road ice cream: Heterogeneous mixture (non-uniform composition containing distinct chunks of nuts, marshmallows, and chocolate).

  3. A bowl of vegetable soup: Heterogeneous mixture (contains distinct visible components such as vegetables and liquid broth).

  4. Mouthwash: Homogeneous mixture (uniform liquid composition throughout).

  5. An unopened can of cola: Homogeneous mixture (uniform liquid solution under pressure).

  6. A dinner salad: Heterogeneous mixture (clearly distinguishable separate ingredients).

States of Matter and Changes of State

• States of Matter: The physical form in which matter exists. The three standard states are solid, liquid, and gas.

  • Solid:

    • Particle Arrangement: Particles are arranged in an orderly, highly structured pattern, tightly packed together, and capable of moving only slightly (vibrating in place).

    • Properties: Has a definite shape and a definite volume.

  • Liquid:

    • Particle Arrangement: Particles are less orderly than in a solid and are able to move freely past one another.

    • Properties: Has a definite volume, but takes the shape of its container.

  • Gas:

    • Particle Arrangement: Particles have no orderly arrangement, are positioned very far apart from one another, move at extremely high speeds, and frequently collide with each other and the interior walls of their container.

    • Properties: Has no definite shape and no definite volume; expands completely to fill whatever container it is placed in.

Physical vs. Chemical Changes

• Physical Change:

  • Occurs when the state or physical form of matter is altered, but its fundamental chemical identity remains unchanged.

  • Examples: Changes of state such as melting ice into liquid water or boiling water into steam.

• Chemical Change:

  • Results in an actual transformation of the chemical identity of a substance, converting it into one or more entirely new substances.

  • When a substance undergoes a chemical change, it is referred to as undergoing a chemical reaction.

Energy, Heat, and Conservation Laws

• Definition of Energy: Energy is defined as the capacity to do work and to transfer heat. Anytime matter is altered or transformed in any way, work has been performed.

• Forms of Energy:

  • Potential Energy: Stored energy held by an object due to its position or chemical structure.

    • Example: Water held stationary behind a hydro-electric dam.

  • Kinetic Energy: The energy of motion.

    • Example: When the dam's floodgates are opened and water pours through, stored potential energy is directly converted into kinetic energy.

• Law of Conservation of Energy: Energy takes various forms, but it is never created and never destroyed; it can only be transformed from one form into another.

• Energy Units:

  • Joule ($\text{J}$): The official standard SI unit for measuring energy.

  • calorie ($\text{cal}$): The quantity of energy needed to raise the temperature of exactly $1$ gram ($\text{g}$) of liquid water by $1$ degree Celsius ($1\,^\circ\text{C}$).

    • Conversion Factor: $1\,\text{calorie} = 4.184\,\text{joules}$

  • Nutritional Calorie ($\text{Cal}$): Used in food labeling, which is $1000$ times larger than a standard scientific calorie.

    • Conversion Factor: $1\,\text{Calorie} = 1000\,\text{calories}$

Practice Problems: Potential vs. Kinetic Energy

• A mountain climber sits at the top of a peak: Mainly Potential Energy.

• A mountain climber rappels down a cliff: Mainly Kinetic Energy.

• A hamburger sits on a plate: Mainly Potential Energy (stored chemical energy).

• A nurse inflates a blood pressure cuff: Mainly Kinetic Energy.

Systems of Measurement and Units

• Structure of Measurements: Every measurement consists of two indispensable parts: a numerical value and a unit. A defined, standardized system of measurement is necessary to compare experimental quantities easily.

• The International System of Units (Système International d’Unités - SI): The modern international standard version of the metric system.

  • Standard SI Unit for Mass: Kilogram ($\text{kg}$).

  • Standard SI Unit for Volume: Liter ($\text{L}$).

  • Standard SI Unit for Length: Meter ($\text{m}$).

• Metric Prefixes: Attached to base SI or metric units to alter the magnitude of the unit by specific powers of 10.

• Customary Systems in Healthcare: Healthcare professionals primarily use SI and metric units, but must remain fluent in conversions involving the U.S. customary system of measurement.

Specific Quantities: Mass and Volume

• Mass vs. Weight:

  • Mass: A quantitative measure of the total amount of material contained within an object. Common laboratory unit is the gram ($\text{g}$).

  • Weight: Determined by the local gravitational pull exerted upon the object, which varies based on geographical location.

  • Earth Surface Equivalency: As long as an object is weighed in roughly the same location on Earth's surface, its measured mass and weight yield the exact same numerical value.

• Volume:

  • A three-dimensional quantitative measure of the physical space occupied by matter.

  • Laboratory Settings: Routinely measured using a graduated cylinder or a pipet. The standard laboratory unit is the milliliter ($\text{mL}$).

  • Clinical Settings: Routinely measured using calibrated syringes. The standard clinical unit is the cubic centimeter ($\text{cc}$ or $\text{cm}^3$).

  • Volume Equivalencies: $1\,\text{mL} = 1\,\text{cm}^3$ and $5\,\text{mL} = 1\,\text{teaspoon}$.

Practice Problems: Unit Magnitude Comparisons

• Which is larger?

  1. $1\,\text{g}$ or $1\,\text{lb}$: $1\,\text{lb}$

  2. $1\,\text{lb}$ or $1\,\text{kg}$: $1\,\text{kg}$

  3. $1\,\text{L}$ or $1\,\text{qt}$: $1\,\text{L}$

  4. $1\,\text{cal}$ or $1\,\text{J}$: $1\,\text{cal}$

Scientific Notation, Accuracy, and Precision

• Scientific Notation Standard Form:

  • Expressed as $C \times 10^n$

  • $C$ is the coefficient, which must be a number greater than or equal to $1$ but strictly less than $10$ (1 \le C < 10).

  • $n$ is an integer exponent indicating the number of tens places (power of 10) applied.

  • Positive Exponent ($+n$): Indicates the actual number is greater than $1$, representing how many times the coefficient must be multiplied by $10$.

  • Negative Exponent ($-n$): Indicates the actual number is between $0$ and $1$, representing how many times the coefficient must be divided by $10$.

  • Significant Figures: In scientific notation, only significant figures are written in the coefficient.

Practice Problems: Scientific Notation Conversions

• $0.0144 = 1.44 \times 10^{-2}$

• $144 = 1.44 \times 10^2$

• $36.32 = 3.632 \times 10^1$

• $0.0000098 = 9.8 \times 10^{-6}$

• Clinical Application Thought Problem: One millionth of a liter of blood ($10^{-6}\,\text{L}$ or $1\,\mu\text{L}$) contains approximately $5$ million red blood cells ($5 \times 10^6\,\text{cells}$).

Accuracy versus Precision

• Accuracy:

  • Refers to how close a measured value is to the actual or true value.

  • Example: If a patient's true body temperature is $103\,^\circ\text{F}$, and two sequential readings yield $98\,^\circ\text{F}$ and $102\,^\circ\text{F}$, the second reading ($102\,^\circ\text{F}$) is significantly more accurate.

• Precision:

  • A measure of numerical reproducibility, representing how close a series of measurements are to one another, regardless of whether they are close to the true value.

  • Example: Evaluating two sets of temperature measurements:

    • Set 1: $98\,^\circ\text{F}$, $99\,^\circ\text{F}$, $97\,^\circ\text{F}$

    • Set 2: $90\,^\circ\text{F}$, $100\,^\circ\text{F}$, $96\,^\circ\text{F}$

    • Set 1 displays tight internal agreement, making it far more precise than Set 2.

• Best Measurement Practice: High-quality scientific data requires both high accuracy and high precision. This is systematically accomplished by taking multiple measurement trials and calculating their mathematical average.

Significant Figures, Exact Numbers, and Rounding Rules

• Definition of Significant Figures: In any physical measurement, the significant figures include all digits known with complete certainty plus exactly one final estimated digit.

• Instrument Precision:

  • Non-digital Devices: The user must estimate the final digit, introducing inherent measurement uncertainty.

  • Digital Devices: Automatically display the correct final count of significant figures on the digital screen.

• Rule for Non-Zero Digits: All non-zero digits ($1$ through $9$) in a measured number are automatically significant.

• Rules for Zeros: Zeros are context-dependent:

  • Zeros without a Decimal Point: If a zero at the end of a large number lacking a decimal point is intended to be significant, it can be denoted by placing an explicit decimal point directly after that zero (e.g., $100.$) or by drawing a horizontal line directly above the significant zero.

• Exact Numbers:

  • Numbers derived from defined conversion factors (e.g., $1\,\text{dL} = 0.1\,\text{L}$) or direct counting of discrete items contain an infinite number of significant figures and do not limit calculations.

Practice Problems: Identifying Significant Figures

• $30.1\,^\circ\text{C}$: Contains $3$ significant figures.

• $0.00730\,\text{m}$: Contains $3$ significant figures (leading zeros are not significant; trailing zero after decimal is significant).

• $7.30 \times 10^3\,\text{m}$: Contains $3$ significant figures.

• $44.50\,\text{mL}$: Contains $4$ significant figures.

Practice Problems: Writing in Exponential Notation

• $7032\,\text{cal} = 7.032 \times 10^3\,\text{cal}$ ($4$ significant figures)

• $88.0\,\text{L} = 8.80 \times 10^1\,\text{L}$ ($3$ significant figures)

• $0.00005\,\text{g} = 5 \times 10^{-5}\,\text{g}$ ($1$ significant figure)

• $0.06430\,\text{lb} = 6.430 \times 10^{-2}\,\text{lb}$ ($4$ significant figures)

Mathematical Operations and Rounding Rules

• Fundamental Constraint: Manipulating measured quantities via arithmetic cannot increase their certainty. An answer can be no more certain than the least certain measurement in the calculation.

• Addition and Subtraction Rule: The calculated result must be rounded to match the least number of decimal places present in any of the original measured numbers.

• Multiplication and Division Rule: The calculated result must be rounded to match the least total number of significant figures present in any of the original measured numbers.

  • Example Question: If a person reports their age as $27$ years, how many days old are they? Calculation must reflect proper precision constraints.

• Standard Rounding Procedure:

  1. Look at the leftmost digit to be dropped (the digit immediately following your designated final retained significant figure).

  2. If that digit is $4$ or less, drop it and all subsequent digits without changing the retained figures.

  3. If that digit is $5$ or greater, increase the final retained digit by $1$ and remove all dropped digits.

  4. Large Numbers Without Decimals: When rounding large numbers that lack decimal points, substitute placeholder zeros for any dropped non-significant digits.

  5. Multi-step Calculations: Do not round intermediate steps. Round only at the very final step to avoid severe cumulative rounding errors.

Practice Problems: Calculations with Significant Figures

• $0.12 \times 1.77 = 0.2124 \rightarrow 0.21$ (rounded to $2$ significant figures)

• $690.4 \div 12 = 57.533... \rightarrow 58$ (rounded to $2$ significant figures)

• $5.444 - 0.44 = 5.004 \rightarrow 5.00$ (rounded to $2$ decimal places)

• $16.5 + 0.114 + 3.55 = 20.164 \rightarrow 20.2$ (rounded to $1$ decimal place)

Percentages and Clinical Applications

• Definition of Percent: Represented by the symbol $\%$, percent literally translates to "part out of 100 total" or hundredths. It provides a standardized metric to compare directly two sets of data that possess different total sample sizes.

• Mathematical Conversions:

  • Fraction to Percent: Divide the numerator by the denominator, multiply the resulting quotient by $100$, and attach a $\%$ symbol.

  • Decimal to Percent: Multiply the decimal value by $100$ and attach a $\%$ symbol.

Clinical and Healthcare Applications of Percent

• Percent Active Ingredient: Used because of the extreme potency of many therapeutic medications; inert chemical binders are added to increase the physical bulk size of a pill to make it handleable.

• Percent of an Adult Dose: Because pediatric patients weigh significantly less than adults, children are commonly prescribed a calculated percentage of the standard adult dose.

• Percent in Nutrition Labeling: Nutritional facts panels list absolute quantities of carbohydrates, proteins, and fats alongside the Percent Daily Value ($\%\text{DV}$), indicating how much a single food serving contributes toward total suggested daily dietary requirements.

Density and Specific Gravity

• Density Formula: Density ($d$) is an intrinsic physical property defined as the ratio of a substance's mass ($m$) to its volume ($V$):     $d = \frac{m}{V}$

• Density of Liquid Water: Exactly $1.00\,\text{g}$ of pure water occupies a volume of $1.00\,\text{mL}$; therefore, the density of water is $1.00\,\text{g/mL}$.

• Density as a Conversion Factor: Because density remains constant for a specific substance at a given temperature, it serves as a conversion factor between mass and volume.

  • Example (Isopropyl Alcohol): At $20\,^\circ\text{C}$, rubbing alcohol has a density of $0.785\,\text{g/mL}$. Thus, $0.785\,\text{g}$ occupies $1.00\,\text{mL}$.

  • Derived Conversion Factors:         $\frac{0.785\,\text{g}}{1.00\,\text{mL}} \quad \text{and} \quad \frac{1.00\,\text{mL}}{0.785\,\text{g}}$

Specific Gravity

• Definition: Specific gravity ($\text{sp\,gr}$) is the dimensionless ratio of the density of a target sample to the density of pure water measured at the same reference temperature:     $\text{sp\,gr} = \frac{\text{density of sample}}{\text{density of water}}$

• Properties: Specific gravity is completely unitless because the mass and volume units in the numerator and denominator cancel out. Water density is $1.00\,\text{g/mL}$ at $4\,^\circ\text{C}$ and remains extremely close to this value at body temperature.

• Measurement Instrument: Liquid specific gravity is measured using a specialized optical instrument called a refractometer.

• Practical Applications: Used to measure battery acid strength in automotive batteries, radiator antifreeze concentration, alcohol percentage during brewing/winemaking, and clinical urine concentration.

Clinical Diagnostic Significance of Urine Specific Gravity

• Definition of Urine: A complex liquid mixture of water and metabolic waste products excreted by the kidneys.

• High Urine Specific Gravity:

  • Indicates an excessively high concentration of dissolved metabolic waste relative to water.

  • Diagnostic of clinical dehydration or pathologically elevated secretion of antidiuretic hormone (ADH), which forces blood serum water retention.

  • High ADH levels are frequently triggered by severe physical stress, body trauma, or recent major surgery.

• Low Urine Specific Gravity:

  • Indicates overly dilute urine.

  • Diagnostic of underlying kidney disease, excessive fluid intake, or severe underproduction of ADH.

Temperature Scales, Thermoregulation, and Specific Heat

• Temperature Definition: A physical measurement of the hotness or coldness of a substance, acquired via a physical thermometer or electronic temperature probe.

• Temperature Scales:

  • Fahrenheit Scale ($^\circ\text{F}$): Standard scale used predominantly in the United States.

  • Celsius Scale ($^\circ\text{C}$): Metric scale used throughout the global scientific community and rest of the world.

  • Kelvin Scale ($\text{K}$): The absolute temperature scale and official SI unit for temperature. The Kelvin and Celsius scales use degree increments of the exact same size, but their zero points are offset by $273$ units ($\text{K} = ^\circ\text{C} + 273$).

  • Scale Conversion Relationship: One degree Celsius is equivalent to $1.8$ degrees Fahrenheit, with their zero points offset by $32$ degrees ($^\circ\text{F} = 1.8(^\circ\text{C}) + 32$).

• Practice Problem: Convert a summer day temperature of $85\,^\circ\text{F}$ to Celsius:     $^\circ\text{C} = \frac{85 - 32}{1.8} = 29.4\,^\circ\text{C}$

Human Body Temperature and Clinical Pathologies

• Normal Human Body Temperature: $98.6\,^\circ\text{F}$ or $37.0\,^\circ\text{C}$ (varies physiologically among individuals and throughout daily circadian rhythms).

• Hyperthermia:

  • Occurs when human body temperature rises above $40.0\,^\circ\text{C}$ ($104\,^\circ\text{F}$).

  • Pathological Effects: Induces severe bodily convulsions, deep coma, or permanent brain damage.

• Hypothermia:

  • Occurs when human body temperature drops below $35\,^\circ\text{C}$ ($95\,^\circ\text{F}$).

  • Pathological Effects: Patient feels intensely cold, develops an irregular heartbeat (arrhythmia), and exhibits a abnormally slow breathing rate.

Heat and Specific Heat Capacity

• Heat Definition: Thermal kinetic energy naturally flowing from a body of higher temperature to a body of lower temperature.

• Specific Heat Capacity (Specific Heat): The exact quantity of heat energy required to raise the temperature of exactly $1\,\text{g}$ of a substance by $1\,^\circ\text{C}$.

  • Metals: Possess characteristically low specific heat values (heat up and cool down rapidly).

  • Water: Possesses a remarkably high specific heat value, enabling substantial thermal regulation.

Dimensional Analysis and Unit Conversions

• Dimensional Analysis Concept: A systematic mathematical problem-solving method utilizing conversion factors to transform a quantity expressed in one unit into an equivalent quantity expressed in a different unit.

• Equivalencies: Physical quantities linked by an equal sign form equivalent units (e.g., $1\,\text{dL} = 0.1\,\text{L}$).

• Step-by-Step Dimensional Analysis Procedure:

  1. Step 1: Identify and write down the final desired unit(s) for the answer.

  2. Step 2: Establish and state all given initial information and starting values.

  3. Step 3: Determine the appropriate conversion factor ratios needed to cancel out unwanted units, placing the desired target unit in the numerator.

  4. Step 4: Perform mathematical multiplication/division and round to the correct significant figures.

• Introductory Example: How many eggs are present in $2$ dozen?     $2\,\text{dozen} \times \frac{12\,\text{eggs}}{1\,\text{dozen}} = 24\,\text{eggs}$

Practice Problems: Clinical Unit Conversions

1. Weight-Based Drug Dosing: A patient weighs $185\,\text{lb}$. Convert this weight into kilograms ($\text{kg}$) to calculate a drug prescription recommended in $\text{mg/kg}$.     $185\,\text{lb} \times \frac{1\,\text{kg}}{2.205\,\text{lb}} = 83.9\,\text{kg}$

2. Serum Volume: A blood serum sample vial contains a volume of $15\,\text{mL}$. Convert this volume into liters ($\text{L}$).     $15\,\text{mL} \times \frac{1\,\text{L}}{1000\,\text{mL}} = 0.015\,\text{L}$

3. Cardiac Stroke Volume: The average volume of blood pumped by a single heart beat is $0.070\,\text{L}$. Convert this volume from liters ($\text{L}$) into customary cups.