Grade 9 Science, All Quarters.

Grade 9 SCIENCE PEPT 1-4th quarter(s) all lessons.

The respiratory system

helps us breathe in oxygen and breathe out carbon dioxide, which is a waste gas.

Parts that are protected and covered by the rib cage.

Nasal cavity, pharynx, larynx, trachea, bronchi, bronchioles, aveoli

At the end of each bronchiole are tiny sacs called

alveoli (singular: alveolus)

Tiny, grape-like clusters surrounded by capillaries

alveoli

A dome-shaped muscle below the lungs

The diaphragm

The fluid part that carries nutrients, hormones, and waste

Plasma

Help in blood clotting

Platelets

Carries blood to tissues and back to the heart

Blood Vessels

Arteries

Carries oxygen-rich blood away from the heart; thick walls to withstand high pressure

Veins

Carry oxygen-poor blood toward the heart; have valves to prevent backflow

Capillaries

Smallest and most numerous; sites of exchange of oxygen, nutrients, and waste with tissue cells in capillary beds

Arterioles

Branch from arteries to capillaries

Venules

Connect capillaries to veins

Pulmonary Circuits

Circulates blood through the lungs for gas exchange (oxygen in, carbon dioxide out)

Systemic Circuits

Circulates blood to all parts of the body to deliver oxygen and nutrients and remove wastes

Left ventricle

The largest chamber because it pumps blood to the entire body. (The right side is smaller as it only pumps to the lungs.)

Heartbeat rate

• Normal resting rate: 60-80 beats per minute

• During exercise: around 120 beats per minute

DNA

 made of nucleotides, the building blocks.

nucleotide consists of three part

1. Phosphate group (P)

2. A 5-carbon sugar (called pentose sugar)

3. An organic nitrogenous base (which gives the nucleotide its identity)

 Four types of nitrogen bases are

• Adenine (A)

• Thymine (T)

• Guanine (G)

• Cytosine (C)

Base pairing rules:

• A pairs with T

• C pairs with G

Tiny structures that you cannot see with the naked eye, They are found on chromosomes, which are long, spaghetti-like structures inside cells.

Genes

Found inside cells, specifically within the nucleus (the cell’s “brain”) are made of DNA and contain many genes (hundreds or thousands)

Chromosomes

 Small, egg-shaped structure that controls cell activities. Contains more information than a dictionary.

Nucleus

Either allele is fully dominant; the offspring show a blend of parental traits (phenotype).

Incomplete Dominance

Both alleles are expressed simultaneously in heterozygotes.

Codominance

More than two alleles exist for some genes, leading to multiple possible phenotypes.

Multiple Alleles

Major cause of wildlife disappearance due to destruction of natural habitats. Causes include:

Illegal logging, Kaingin (slash-and-burn farming), Forest fires, Typhoons, Industrialization and land conversion

Deforestation

Eutrophication (excess nutrients from sewage, manure, pesticides, garbage, carcasses) leading to water quality decline.

Water Pollution

Caused by harmful chemicals from human activities:

Transportation

Industry and mining

Manufacturing and vehicle combustion.

Air Pollution

Overexploitation and overfishing

Dynamite fishing and muro-ami (blast fishing)

Deforestation, agriculture, and mining activities

Coastal development—hotels, roads, resorts, and housing.

Destruction of Coastal Resources

Includes any form of precipitation (rain, snow, sleet, fog) that contains acidic components, such as sulfuric acid or nitric acid. It results from emissions of sulfur oxides (SOx) and nitrogen oxides (NOx) from factories and fuel combustion.

Acid Precipitation (Acid Rain)

A process by which autotrophic organisms (like plants) make food using sunlight, carbon dioxide (CO₂), and water. Produces oxygen and glucose (a simple carbohydrate that stores energy).

Photosynthesis

Surrounded by double membranes.

Contains stacked structures called thylakoids (where light reactions occur)

Chloroplast

Embedded with chlorophyll (pigment).

Light-absorbing molecules that start photosynthesis

Thylakoids

Takes place in the thylakoid membranes in the presence of light.

Converts light energy into chemical energy.

Involves photosystems I and II (complexes of pigments).

Chlorophyll absorbs violet, blue, and red light, reflecting green.

Water splits (oxidation), releasing oxygen and electrons.

Energy boosts electrons to a higher state (photosynthetic electron transport chain).

Light-dependent Reaction

Occurs in the stroma of chloroplasts.

Uses chemical energy from light reactions to produce glucose from CO₂.

Calvin Cycle (Dark Reaction)

Gas Exchange

• Surrounding guard cells control opening and closing based on osmotic changes.

• Regulate the intake of CO₂ and release of O₂.

Stomata

Photosynthesis

stores solar energy in glucose.

Cellular Respiration

Occurs in three main stages: Glycolysis, Krebs Cycle, and Electron Transport Chain.

Breaks down glucose to produce ATP, the main energy currency of cells.

Glycolysis

Located in the Cytoplasm.

Process:

Glucose (6-carbon) splits into 2 pyruvate molecules.

Uses 2 ATP, produces 4 ATP (net 2 ATP).

Generates 2 NADH molecules for energy transfer.

Krebs Cycle (Citric Acid Cycle)

Located in the Mitochondrial matrix.

Process:

Pyruvate converts to acetyl-CoA, which enters the cycle.

Produces NADH, FADH₂, and ATP. Releases CO₂ as waste.

Restores oxaloacetate for the next cycle.

Electron Transport Chain

Located in the Inner mitochondrial membrane.

Process:

NADH and FADH₂ donate electrons to transport chains.

Electrons move along, pumping H+ ions to create a chemiosmotic gradient.

ATP synthase uses this gradient to produce ATP.

The electrons finally combine with oxygen to form water.

Electron Configuration

The arrangement of electrons around the nucleus of an atom. Helps predict element properties and chemical behavior. Format: Number + Letter + Superscript
Number: Principal energy level (shell)

Letter: Sublevel (s, p, d, f)

Superscript: Number of electrons in that sublevel

Filling Electron Subshells

• Electrons fill lowest energy orbitals first (Aufbau Principle).

  The order of filling:

• 1s→2s→2p→3s→3p→4s→3d→4p→5s→4d→5p→6s→4f→5d→6p→7s→5f→6d→7p

• 1s→2s→2p→3s→3p→4s→3d→4p→5s→4d→5p→6s→4f→5d→6p→7s→5f→6d→7p

Aufbau Principle

• Electrons occupy orbitals starting with the lowest energy levels.

Pauli Exclusion Principle

An orbital can hold a maximum of 2 electrons, with opposite spins.

Hund’s Rule

When electrons occupy degenerate orbitals (same energy), they fill each one singly first with parallel spins, then pair up.

Short Electron Configurations

Use noble gas abbreviation: replace the configuration of a noble gas with its symbol in brackets. Example: [Ar]4s23d6

Quantum Numbers

 they specify the possible locations and states of electrons in an atom. Has a unique set of four quantum numbers: n, ℓ, mℓ, ms

Principal Quantum Number (n)

• Represents the main energy level or shell.

• Values: 1, 2, 3, 4, ...

• Distance from nucleus:

  • n=1

  • n=1 = closest to nucleus

  • n=2

  • n=2 = farther out, and so on.

• Also corresponds to the row number in the periodic table.

Angular Momentum Quantum Number (ℓ)

Describes the shape of the orbital. Values depend on n

Range: 0 to 

n−1

n−1.

ℓ=0

ℓ=0 = s orbital (spherical)

ℓ=1

ℓ=1 = p orbital (dumbbell-shaped)

ℓ=2

ℓ=2 = d orbital (cloverleaf)

ℓ=3 ℓ=3 = f orbital (complex shape).

Magnetic Quantum Number (mℓ)

Describes the orientation of the orbital in space. Values: −ℓ to +ℓ (integers).

Spin Quantum Number (ms)

Describes the electron’s spin.

Possible values: +1/2 or −1/2.

Electrons in the same orbital must have opposite spins (paired spins).

Spin is represented as arrows: ↑ (+1/2) or ↓ (−1/2).

Elements

They combine to form compounds held together by strong chemical bonds.

Ionic and covalent.

Two main types of bonds

Ionic Bonds

Formed when a metal transfers electrons to a non-metal.

Cation

Positively charged ion (lost electrons).

Anion

Negatively charged ion (gained electrons).

Covalent Bonds

Formed when two non-metals share electrons.

Neither atom fully gains or loses electrons; they share to complete their octet.

Electronegativity

Atom's ability to attract electrons. Increases from left to right across periods and top to bottom within groups. Metals: low electronegativity, tend to lose electrons. Non-metals: high electronegativity, tend to gain electrons

Ionic bond

difference greater than 1.9

Chemical Compounds

Formed when two or more different atoms chemically bond.

Cannot be separated easily.

Types:

Ionic Compounds: formed by metal and non-metal transfer of electrons.

Covalent Compounds: formed by sharing electrons between non-metals.

Ionic Compounds

• Formation: Metal reacts with non-metal, resulting in ions (cations and anions).

• Physical Properties:

  1. High melting and boiling points due to strong electrostatic attraction between ions. Usually solids, hard but brittle; can break under force. High electrical and thermal conductivity when dissolved in water; conduct electricity as liquid ions. Insulators in solid form — do not conduct electricity as solids.

Covalent (Molecular) Compounds

Formation: Non-metals share electrons to attain a stable octet.

• Physical Properties:

  1. Low melting and boiling points; easily separated.

  2. Usually soft, brittle, and can be solids, liquids, or gases.

  3. Generally insoluble in water.

• Examples: Carbon dioxide (CO₂), water, ammonia.

An atom

is neutral when number of protons = electrons.

An ion

is formed when electrons are gained or lost, resulting in an unequal number of protons and electrons. it also carries an electric charge. (cation, anion)

Valence electrons

Outermost electrons responsible for bonding.

Protons

 It cannot be gained or lost — changing them would change the element.

Monatomic

Single atom with a single charge.

Polyatomic

group of atoms with a collective charge.

Ions

 vital for body functions. They maintain osmotic pressure.

• Enable muscle and nerve activity.

• Create phenomena like auroras when charged particles interact with Earth's atmosphere

Period

Horizontal row, indicates the number of valence shells.

Family (Group)

Vertical column, indicates the number of valence electrons.

Lewis Dot Structures (LEDS)

• Graphical way to show valence electrons.

  Symbols: element symbol with dots representing valence electrons.

  • Dots are placed clockwise around the symbol.

  • Helps visualize bond formation and lone pairs

Chemical Bonds

are attraction forces that link atoms together. Formed when valence electrons interact.

Covalent Bonds

• Occur when electrons are shared.

• Equal or similar electronegativity (difference ≤ 0.4):

  • Nonpolar covalent bond (e.g., F₂): electrons shared equally.

• Moderate difference (0.4 < difference < 1.9):

  • Polar covalent bond (e.g., H₂O): electrons shared unequally, resulting in δ+ and δ- charges.

Ionic Bonds

• Occur when electrons are fully transferred.

• Electronegativity difference > 1.9:

  • Metal transfers electrons to a non-metal.

  • Results in ions:

    • Cations: positive charge (lose electrons).

    • Anions: negative charge (gain electrons).

  • Form ionic compounds (e.g., NaCl).

Carbon

exists in various forms such as diamonds, graphite, and organic compounds.

It is crucial for life on Earth, providing energy, forming the backbone of biological molecules, and influencing climate regulation.

Unique bonding ability:

Contains 4 valence electrons.

Forms covalent bonds by sharing electrons.

Capable of forming long chains, branched structures, and cyclic molecules.

Can form single, double, or triple bonds.

Carbon compounds exhibit isomerism:

molecules with the same molecular formula but different structures.

Found in non-living environments (atmosphere, rocks, fossil fuels) and living organisms (proteins, carbohydrates, lipids, nucleic acids).

Can form molecules ranging from simple two-atom molecules like carbon monoxide (CO) to complex structures like DNA.

Hydrocarbons

Organic compounds made up exclusively of carbon and hydrogen atoms. They are primarily classified into three types based on their bonding

Alkanes

Contain only single bonds between carbon atoms.

Also called paraffins.

Saturated hydrocarbons.

General formula: CₙH₂ₙ+2.

Suffix: -ane (e.g., methane, ethane, propane).

Structure: Carbon atoms bonded to four other atoms.

Alkenes

Contain one or more double bonds.

Also called olefins.

Unsaturated hydrocarbons.

General formula: CₙH₂ₙ.

Suffix: -ene.

Example: Ethene.

Alkynes

Contain one or more triple bonds.

Also called acetylenes.

General formula: CₙH₂ₙ−2.

Suffix: -yne.

Example: Ethyne.

Alcohols

Contain the -OH group.

Used as disinfectants, solvents, and fuels.

Naming: locate the longest chain with -OH, add suffix -ol.

Example: Butanol (from butane).

Ethers

Contain an oxygen atom bonded to two alkyl groups (R-O-R).

Used as solvents; extremely flammable.

Naming: specify alkyl groups alphabetically + ether.

Example: Dimethyl ether.

Aldehydes

Contain a carbonyl group (C=O) at the end of a chain.

Generic formula: R-CHO.

Naming: suffix -al.

Example: Ethanal.

Ketones

Contain a carbonyl group within the chain (C=O connected to two carbons).

Naming: suffix -anone, with position number.

Example: 2-Propanone (acetone)

Carboxylic Acids

Contain -COOH group.

Example: Butanoic acid (from butane).

Esters

Derived from carboxylic acids and alcohols (RCOOR).

Used for pleasant odors and flavors.

Naming: first alkyl, then stem of acid + -oate.

Example: Methyl butanoate.

Amines

Contain basic nitrogen atoms (NH₂, NH, N).

Derived from ammonia (NH₃).

Naming: add -amine to the alkane part.

Example: Methylamine.

Amides

Derived from carboxylic acids and amines.

Used as drugs (e.g., penicillin).

Naming: replace -ic acid with -amide.

Example: Acetamide.

Atomic mass

Average mass of an atom of an element.

Molecular mass

Average mass of a molecule; sum of atomic masses in a molecule.

Formula mass

 Sum of atomic masses in a formula unit of an ionic compound; applicable to ionic substances.

SI unit for atomic/molecular mass

atomic mass unit (amu)

SI unit for molar mass

grams per mole (g/mol)

The Mole (mol)

The amount of substance containing.
Chemists relate atoms and molecules to their masses using it as a unit of measurement.

Atomic mass

expressed in grams of an element.

Molar mass

expressed in grams per mole (g/mol) for compounds or elements.