Chapter 1-9 Review Flashcards (Matter, Measurements, and Periodic Table Basics)
Matter and States of Matter
Matter definition: anything that occupies space; examples include a cell phone and the human body.
Three common states of matter: solid, liquid, gas.
Solid: has a definite shape (e.g., ice).
Liquid: takes the shape of its container.
Gas: has no fixed shape.
Phase transition example: solid → liquid → gas (ice → water → water vapor).
Visual of states: solid molecules are tightly packed; liquid molecules are more spread out; gas particles are far apart.
Brief note on plasma: not covered in-depth here; focus is on solid, liquid, gas.
Mass vs Weight
Mass: a measure of the amount of matter in an object (how much matter is present; it occupies space).
Weight: the gravitational force exerted on that mass; depends on gravity.
On Earth, mass and weight are related via gravity; on the Moon, gravity is weaker, so weight decreases, but mass remains.
Example understanding:
An object has mass whether or not gravity is acting on it; weight is what a scale measures under gravity.
Quick takeaway:
Mass = amount of matter (intrinsic).
Weight = gravitational force on that mass (external).
Reiteration to solidify concept:
On Earth, you weigh something on a scale; on the Moon you would weigh less due to lower gravity, but the object’s mass is unchanged.
Conservation of Matter
Law of conservation of matter: there is no detectable change in the total quantity of matter when it is converted from one form to another (e.g., solid ↔ liquid ↔ gas).
Practical implication: if you start with 10 molecules in a solid and melt them, you still have 10 molecules in the liquid (the arrangement changes, not the count).
Analogy/connection: relates to the broader principle that energy can transform rather than be created or destroyed; similar to “energy is conserved” ideas in chemistry and physics.
Mixtures
Homogeneous mixture: uniform composition; appears the same throughout (e.g., Kool-Aid in water).
Heterogeneous mixture: non-uniform; contains visibly different components (e.g., oil and vinegar in salad dressing).
Physical vs Chemical Changes (Properties and Changes)
Physical change: reversible with the right conditions; the original substance can be recovered.
Example sequence: solid ice → liquid water → steam (gas); reverse by cooling/condensing.
Chemical change: irreversible in the sense of returning to the original substance using ordinary means; composition changes (new substances formed).
Examples given: rusting of metal; chemical relaxers for hair (permanent straightening).
Take-home message: physical changes can be reversed to regain the original material; chemical changes result in a new substance with different properties.
Measurements and Standard Units (Conversions)
Standardization goal: scientists use common base units to read and compare data globally.
Base SI units introduced:
Length: meter (m)
Mass: kilogram (kg)
Time: second (s)
Temperature: Kelvin (K) (note: Celsius and Fahrenheit are still used, but Kelvin is the SI standard for science)
Temperature conversation norms referenced:
Kelvin = Celsius + 273.15, i.e., K = °C + 273.15
Common conversion rules highlighted (the exact conversion factors are given and must not be changed):
Length: 1 inch = 2.54\,\text{cm}
Mass: 1 kilogram = 2.2\,\text{lb}
Why standardization matters: allows consistent interpretation of data across countries and disciplines; avoids needing to manually convert in every document.
Step-by-step conversion technique emphasized (dimensional analysis / cancellation):
Put the known value with its unit, add the conversion factor with the unit you want in the opposite position to cancel the old unit and leave the desired unit.
Example approach highlighted with a purple box to emphasize the immutable conversion law.
Worked examples (as demonstrated):
Convert 12 inches to centimeters:
Setup: 12\,\text{in} \times \frac{2.54\,\text{cm}}{1\,\text{in}} = 30.48\,\text{cm}
Convert 45 centimeters to inches:
Setup: 45\,\text{cm} \times \frac{1\,\text{in}}{2.54\,\text{cm}} = 17.72\,\text{in}
Convert 62 pounds to kilograms:
Setup: 62\,\text{lb} \times \frac{1\,\text{kg}}{2.2\,\text{lb}} \,=\, 28.18\,\text{kg}
Convert 78 kilograms to pounds:
Setup: 78\,\text{kg} \times \frac{2.2\,\text{lb}}{1\,\text{kg}} \,=\, 171.6\,\text{lb} (rounded to 172 lb in practice)
Note: the method scales to multi-step conversions (stoichiometry in chemistry contexts).
Practical note on dose safety and pharmacology:
Accurate unit conversion is critical when calculating medication dosages for different patient groups; incorrect math can have life-threatening consequences.
The speaker emphasizes verifying calculations and challenging results when something looks off to prevent harm.
Take-home message on conversions:
The conversion factors stay constant; don’t “invent” new values. Use the official ratios (e.g., 1 inch = 2.54 cm; 1 kg = 2.2 lb).
Quick temperature reference: conversions between Fahrenheit, Celsius, and Kelvin exist, with Kelvin usage as the standardized scientific unit; the main explicit relation shown is K = °C + 273.15.
The Periodic Table, Elements, Atoms, and Compounds
Historical context: Dmitri Mendeleev organized elements into a periodic table by weight and properties; he predicted missing elements and gaps in the table.
Key definitions:
Element: a substance that cannot be broken down into simpler substances by ordinary chemical means.
Compound: two or more elements chemically bonded together (e.g., table salt NaCl; water H2O).
Molecule: a cluster of two or more atoms held together; can be an element (O2) or a compound (H2O). The lecturer’s simplification: molecules are two joined things; compounds are “dry” combinations like salt (NaCl) whereas water (H2O) is a molecule formed from hydrogen and oxygen.
Elementary components of matter:
Atoms are composed of protons (positive), neutrons (neutral) in the nucleus, and electrons (negatively charged) orbiting the nucleus.
Electrons contribute negligibly to atomic mass; the mass comes mainly from protons and neutrons.
Atomic number and mass concepts:
Atomic number Z = number of protons.
Mass number A = Z + N (where N is the number of neutrons).
For a neutral atom, the number of protons equals the number of electrons.
Example overview of specific elements (as used in the lecture):
Oxygen: atomic number Z = 8; mass A ≈ 16; neutrons N = A − Z = 8; electrons = 8 in neutral state; outer shell valence electrons = 6; group associated with valence = 6.
Sulfur: Z = 16; A = 32; N = 16; electrons = 16; outer valence electrons = 6; group 6.
Sodium: Z = 11; A = 23; N = 12; electrons = 11; outer valence = 1; group 1.
Calcium: Z = 20; A = 40; N = 20; electrons = 20; electron configuration shells: 2, 8, 8, 2; outer valence = 2; group 2.
Electron shells and valence concept:
First shell maximum 2 electrons; second shell maximum 8; subsequent shells follow similar capacity rules.
Valence electrons: electrons in the outermost shell; determine chemical behavior and the group placement in simplified teaching (groups 1–8 in the lecture’s scheme).
Group placement example: Oxygen has 6 valence electrons → placed in group 6; Sodium has 1 valence electron → group 1; Calcium has 2 valence electrons → group 2.
Periodic table structure insights covered:
Weight increases as you move down the table (in the traditional arrangement by atomic weight).
Some parts of the table were historically moved (lanthanides/actinides) to a bottom block to simplify the page.
The lecture focuses on the practical idea that outer-shell (valence) electron count governs column (group) placement in this course.
Bonding primer (foreshadowing for next week):
Group 8 (noble gases) are “old money”: they have full valence shells (eight outer electrons) and are largely non-reactive.
Groups 1–7 are “new money”: they tend to gain or share electrons to reach a noble-gas configuration (eight valence electrons in early teaching).
This leads to covalent bonding (sharing electrons) and ionic bonding (transfer of electrons) as ways to achieve stable electron configurations.
Easy-to-remember visual metaphors from the lecture (for engagement):
“Electrons buzz around like flies around a rotting mango.” Use rings to organize electrons into shells.
“Old money” vs “new money” analogy to describe noble gases vs reactive elements; “royalty” (noble gases) vs those seeking stability.
Practical goal for students:
By the end of the course, students should be able to interpret the periodic table in terms of valence electrons and predict group placement and bonding tendencies.
Protons, Neutrons, Electrons (Illustrative Examples)
Visualization with oxygen (as used in the lecture):
Oxygen nucleus: Z = 8 protons; mass ~16 implies N = 8 neutrons; electrons = 8 (P = E for neutral atoms).
Electron arrangement example: 2 electrons in the first shell, 6 in the second shell (total 8).
Valence electrons for oxygen = 6; places it in group 6.
Visualization with sulfur:
Z = 16; A = 32; N = 16; electrons = 16; valence electrons = 6; group 6.
Visualization with sodium:
Z = 11; A = 23; N = 12; electrons = 11; shells: 2, 8, 1; valence electrons = 1; group 1.
Visualization with calcium:
Z = 20; A = 40; N = 20; electrons = 20; shells: 2, 8, 8, 2; valence electrons = 2; group 2.
Important rule: P (protons) = E (electrons) in neutral atoms; Neutrons can vary and contribute to mass but not to charge.
Bonding Concepts (Preview for Next Lessons)
Noble gases (group 8) tend to be chemically inert due to complete valence shells; they are the “royal family” in the metaphor.
Other elements seek electrons to fill their outer shell (aim for 8 valence electrons in this teaching framework): this drives bonding.
Bonding types to be covered next: covalent bonding (sharing electrons) and ionic bonding (transfer of electrons).
Ethical, Philosophical, and Practical Implications
Real-world relevance of chemistry literacy:
Accurate measurements, unit conversions, and understanding of mass vs weight are essential in healthcare, pharmacology, and laboratory work.
Mistakes in calculation can have severe consequences (as illustrated by a prescription error scenario), underscoring the ethical responsibility of healthcare professionals to verify and challenge calculations.
Emphasis on verification and critical thinking:
Do not rely solely on someone else’s math; be prepared to check calculations yourself.
Clear, organized problem-solving processes (like the dimensional analysis method) are essential in professional practice.
Summary Takeaways
Matter exists in three common states (solid, liquid, gas) with distinct structural and shape properties.
Mass vs weight distinction is core: mass is intrinsic; weight depends on gravity.
Matter is conserved during phase changes; the count of particles remains constant though their arrangement changes.
Mixtures can be homogeneous or heterogeneous; physical and chemical changes have different reversibility and outcomes.
SI units standardize measurements; common conversions (inch↔cm, kg↔lb, °C↔K) rely on fixed conversion factors.
The periodic table organizes elements by weight and by their outer electron configuration, informing chemical behavior and bonding potential.
Atomic structure basics: protons, neutrons in the nucleus; electrons orbit and determine chemical behavior; Z equals the number of protons; A equals Z plus N; neutral atoms have P = E.
Bonding concepts set the stage for understanding how atoms achieve stable electron configurations, leading to the vast diversity of substances.
Key Equations and Facts (as presented in the lecture)
Conversion: 1\text{ inch} = 2.54\ \text{cm}
Mass-to-weight reference (typical conversion used): 1\text{ kg} = 2.2\ \text{lb}
Temperature relation: K = °C + 273.15
Atomic structure notes:
Z = \text{number of protons}
A = Z + N,\quad N = A - Z
For a neutral atom: P = E (protons equal electrons)
Example element data (illustrative):
Oxygen: Z=8,\ A\approx 16,\ N=8,\ ext{valence} = 6
Sodium: Z=11,\ A=23,\ N=12,\ ext{valence} = 1
Calcium: Z=20,\ A=40,\ N=20,\ ext{valence} = 2