GACE Middle Grades Science (4–8) (705)

All matter consists of atoms, the fundamental building blocks of chemical substances. An atom contains three subatomic particles: protons, neutrons, and electrons.

Neutrons carry no charge and also reside in the nucleus. Neutrons contribute to atomic mass but not to charge. The number of neutrons can vary within an element, creating isotopes. Carbon-12 has six neutrons while carbon-14 has eight neutrons. Both are carbon because both have six protons.

Electrons carry a negative charge and orbit the nucleus in energy levels called electron shells. In a neutral atom, the number of electrons equals the number of protons. Electrons determine chemical behavior and bonding properties.

The nucleus contains nearly all the mass of an atom. Protons and neutrons each have a mass of approximately one atomic mass unit. Electrons have negligible mass, about 1/1836 the mass of a proton.

Atomic mass equals the sum of protons and neutrons. Mass number is written as a superscript before the element symbol. Carbon-14 is written as ¹⁴C, indicating six protons and eight neutrons.

Electrons occupy energy levels at specific distances from the nucleus. The first energy level holds a maximum of two electrons. The second level holds eight electrons. The third level holds eighteen electrons, though it typically fills to eight before the fourth level begins accepting electrons.

Valence electrons occupy the outermost energy level. These electrons participate in chemical bonding. Elements in the same group of the periodic table have the same number of valence electrons, giving them similar chemical properties.

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, the stable configuration of noble gases. Hydrogen is an exception, seeking two valence electrons to match helium.

Dmitri Mendeleev organized elements by increasing atomic mass and similar properties. The modern periodic table arranges elements by increasing atomic number.

Periods are horizontal rows. Elements in the same period have the same number of electron shells. Period 1 elements have one shell, period 2 elements have two shells, and so forth.

Groups are vertical columns. Elements in the same group have the same number of valence electrons and similar chemical properties. Group 1 elements have one valence electron. Group 18 elements have eight valence electrons (except helium with two).

Metals occupy the left and center portions of the periodic table. Metals are shiny, malleable, ductile, and conduct heat and electricity. Most elements are metals.

Nonmetals occupy the upper right portion. Nonmetals are typically dull, brittle, and poor conductors. Nonmetals include carbon, nitrogen, oxygen, sulfur, and the halogens.

Metalloids fall along the staircase line separating metals from nonmetals. Metalloids have properties intermediate between metals and nonmetals. Silicon and germanium are semiconductors used in electronics.

Alkali metals (Group 1) are highly reactive metals with one valence electron. They react vigorously with water to form hydroxides and hydrogen gas. Reactivity increases down the group.

Alkaline earth metals (Group 2) have two valence electrons. They are reactive but less so than alkali metals. Calcium and magnesium are essential for biological processes.

Halogens (Group 17) have seven valence electrons and need one more to complete their octet. They are highly reactive nonmetals that form salts with metals. Fluorine is the most reactive element.

Noble gases (Group 18) have complete valence shells and are chemically inert under normal conditions. They exist as monatomic gases.

Transition metals occupy groups 3 through 12. They have variable oxidation states and often form colored compounds. Many serve as catalysts in chemical reactions.

Atoms bond to achieve stable electron configurations. Three primary bond types exist: ionic, covalent, and metallic.

Ionic bonds form when electrons transfer from one atom to another. Metals lose electrons to become positive ions (cations). Nonmetals gain electrons to become negative ions (anions). Opposite charges attract, holding the ions together.

Sodium chloride exemplifies ionic bonding. Sodium loses one electron to become Na⁺. Chlorine gains that electron to become Cl⁻. The electrostatic attraction between these ions creates an ionic compound.

Ionic compounds form crystal lattices with alternating positive and negative ions. They have high melting points, conduct electricity when dissolved or melted, and are typically soluble in water.

Covalent bonds form when atoms share electrons. Nonmetals bond covalently with other nonmetals. Shared electrons count toward the valence shell of both atoms.

Single bonds share one pair of electrons. Double bonds share two pairs. Triple bonds share three pairs. Bond strength increases with the number of shared pairs.

Water demonstrates covalent bonding. Each hydrogen atom shares its single electron with oxygen. Oxygen shares one electron with each hydrogen. All three atoms achieve stable configurations.

Polar covalent bonds result when electrons are shared unequally. More electronegative atoms attract shared electrons more strongly. Water is polar because oxygen attracts electrons more than hydrogen does.

Nonpolar covalent bonds share electrons equally. Diatomic molecules of the same element (O₂, N₂, H₂) have nonpolar bonds.

Metallic bonds involve a sea of delocalized electrons shared among all metal atoms in a sample. This electron sea explains why metals conduct electricity and heat, are malleable, and are ductile.

Physical properties can be observed without changing the substance's identity. Color, density, melting point, boiling point, hardness, and conductivity are physical properties.

Chemical properties describe how a substance reacts with other substances. Flammability, reactivity with acids, and oxidation tendency are chemical properties. Observing chemical properties requires a chemical change.

Extensive properties depend on the amount of substance. Mass, volume, and length are extensive properties. More matter means greater values for these properties.

Density equals mass divided by volume. Density is an intensive property useful for identifying substances. Water has a density of 1 g/mL at 4°C. Objects with density less than water float; those with greater density sink.

Matter exists in four common states: solid, liquid, gas, and plasma.

Liquids have definite volume but take the shape of their container. Particles are close together but can slide past each other. Liquids flow and are nearly incompressible.

Gases have neither definite shape nor definite volume. They expand to fill any container. Particles are far apart and move rapidly in random directions. Gases are highly compressible.

Plasma is ionized gas with free electrons and positive ions. Plasma occurs at extremely high temperatures. Stars, lightning, and neon signs contain plasma.

Phase changes occur when matter absorbs or releases energy.

Melting converts solid to liquid. Heat energy breaks some intermolecular attractions, allowing particles to move past each other. The temperature at which this occurs is the melting point.

Freezing converts liquid to solid. Energy is released as particles slow and lock into fixed positions. Freezing point equals melting point for a given substance.

Vaporization converts liquid to gas. Boiling occurs throughout the liquid at the boiling point. Evaporation occurs only at the surface at temperatures below the boiling point.

Condensation converts gas to liquid. Energy is released as gas particles slow and intermolecular attractions increase.

Sublimation converts solid directly to gas without passing through the liquid phase. Dry ice (solid CO₂) sublimates at room temperature.

Deposition converts gas directly to solid. Frost forms by deposition of water vapor.

During phase changes, temperature remains constant even though energy is being added or removed. This energy changes the arrangement of particles rather than their kinetic energy.

Pure substances have constant composition and properties. Elements contain only one type of atom. Compounds contain two or more elements in fixed ratios.

Mixtures contain two or more substances that retain their individual properties. Components can be separated by physical means.

Homogeneous mixtures have uniform composition throughout. Solutions are homogeneous mixtures. Salt water looks the same throughout because salt particles are dispersed evenly.

Heterogeneous mixtures have non-uniform composition. Components can be visually distinguished. Salad, soil, and granite are heterogeneous mixtures.

Solutions consist of solute dissolved in solvent. The solvent is present in greater quantity. Water is called the universal solvent because it dissolves many substances.

Solubility is the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature. Temperature generally increases solubility of solids in liquids but decreases solubility of gases in liquids.

Saturated solutions contain the maximum amount of dissolved solute. Unsaturated solutions can dissolve more solute. Supersaturated solutions contain more than the maximum amount, achieved by careful cooling or evaporation.

Separation techniques exploit differences in physical properties. Filtration separates solids from liquids based on particle size. Distillation separates liquids based on boiling point differences. Chromatography separates substances based on their attraction to a stationary phase versus a mobile phase.

Chemical reactions rearrange atoms to form new substances. Reactants are consumed; products are formed. Chemical equations represent reactions using formulas and symbols.

The law of conservation of mass states that mass is neither created nor destroyed in chemical reactions. The total mass of reactants equals the total mass of products. Atoms are rearranged, not created or destroyed.

Synthesis reactions combine two or more reactants into a single product. The general form is A + B → AB. Metals combining with oxygen to form oxides are synthesis reactions.

Decomposition reactions break a single compound into two or more products. The general form is AB → A + B. Electrolysis of water producing hydrogen and oxygen is decomposition.

Single replacement reactions occur when one element replaces another in a compound. The general form is A + BC → AC + B. More reactive elements displace less reactive elements.

Double replacement reactions involve exchange of ions between two compounds. The general form is AB + CD → AD + CB. Precipitation reactions and neutralization reactions are double replacement reactions.

Combustion reactions involve rapid combination with oxygen, releasing heat and light. Hydrocarbons burn to produce carbon dioxide and water. Complete combustion requires sufficient oxygen.

Exothermic reactions release energy to the surroundings. Products have less energy than reactants. Combustion and neutralization are exothermic.

Endothermic reactions absorb energy from the surroundings. Products have more energy than reactants. Photosynthesis and the decomposition of water are endothermic.

Reaction rates depend on several factors. Higher temperature increases reaction rate by increasing particle collisions. Higher concentration increases collision frequency. Catalysts lower activation energy without being consumed. Greater surface area exposes more reactant particles to collisions.

Acids produce hydrogen ions (H⁺) in water. Acids taste sour, turn blue litmus red, and react with metals to produce hydrogen gas. Common acids include hydrochloric acid, sulfuric acid, and acetic acid (vinegar).

Bases produce hydroxide ions (OH⁻) in water. Bases taste bitter, feel slippery, and turn red litmus blue. Common bases include sodium hydroxide, ammonia, and baking soda.

The pH scale measures hydrogen ion concentration. The scale runs from 0 to 14. Acidic solutions have pH below 7. Neutral solutions have pH of 7. Basic solutions have pH above 7.

Each pH unit represents a tenfold difference in hydrogen ion concentration. A solution with pH 3 has ten times more hydrogen ions than a solution with pH 4.

Neutralization reactions occur when acids and bases combine. The products are water and a salt. HCl + NaOH → NaCl + H₂O represents neutralization.

Indicators change color depending on pH. Litmus paper turns red in acid and blue in base. Phenolphthalein is colorless in acid and pink in base. Universal indicator shows a range of colors corresponding to different pH values.