Water Molecule: Polarity, O-H Bonds & Properties

The water molecule exhibits oxygen-hydrogen (O-H) bonds, which are polar covalent bonds due to the significant difference in electronegativity between oxygen and hydrogen. Oxygen atoms are more electronegative than hydrogen atoms. The electronegativity difference leads to unequal sharing of electrons. Electrons are drawn more towards the oxygen atom. This creates a partial negative charge on the oxygen and partial positive charges on the hydrogens. This charge distribution in the water molecule makes it a polar molecule. Polarity is essential for water’s unique properties. These properties include its ability to act as a versatile solvent and its high surface tension.

Ever wonder why some things mix well, like sugar in your coffee, while others, like oil and water, just refuse to get along? The secret lies in something called chemical polarity! It’s like a tiny tug-of-war happening at the atomic level, and it dictates how molecules behave, interact, and ultimately, how the world around us works.

Think of it this way: atoms are like kids sharing toys (electrons). But some kids (atoms) are greedier than others, pulling the toys (electrons) closer to themselves. This unequal sharing creates a slight imbalance of charge, giving rise to polarity. This difference in electronegativity between atoms forms a chemical bond.

In this post, we’re going to unravel the mysteries of polarity, covering everything from electronegativity (the “greediness” of atoms) to dipole moments (a measure of that imbalance), partial charges, and polar covalent bonds. We will touch on the importance of polarity in the properties of water and alcohols and how it affects hydrogen bonding and intermolecular forces. We’ll also explore how polarity dictates solubility (“like dissolves like”) and even plays a role in acidity and basicity.

To get you hooked, here’s a mind-blowing fact: Polarity is what allows soap to wash away grease! Soap molecules have one end that loves water (polar) and another that loves grease (nonpolar). This allows them to bridge the gap and carry the grease away with the water. Pretty neat, huh? Or even think about why water is so essential for life; it’s all because of its polarity! So, buckle up, because we’re about to dive into the electrifying world of polarity!

Electronegativity: The Tug-of-War for Electrons

Okay, so imagine atoms are tiny little people, and electrons are like their favorite toys. Electronegativity is basically how greedy each atom is for those toys when they’re sharing (or trying to share) in a chemical bond. It’s the power of an atom to hog those negatively charged electrons in a bond. Some atoms are generous and don’t mind sharing, while others are like, “Mine! All mine!”

Now, grab your periodic table – it’s the cheat sheet to understand this electron-grabbing behavior! Electronegativity follows a pretty neat pattern. As you move from left to right across a period (a row), the electronegativity generally increases. Think of it like the atoms are getting stronger and more determined to snatch those electrons. But as you move down a group (a column), electronegativity decreases. Those atoms are getting lazier and less interested in electron-grabbing as you go down. This is because the electrons are getting further away from the positive nucleus.

Decoding the Pauling Scale: Ranking the Electron Hogs

To put a number on this greediness, we use something called the Pauling scale. Linus Pauling, the genius behind it, assigned electronegativity values to all the elements. It’s like a leaderboard for electron-grabbing! For example, oxygen (O) is a notorious electron hog with a high electronegativity value (around 3.44). Nitrogen (N) is also pretty greedy (around 3.04). Carbon (C) is more moderate (around 2.55), and hydrogen (H) is relatively chill (around 2.20). Fluorine, the most electronegative element, tops the scale.

The Electronegativity Difference: A Bond’s Destiny

The difference in electronegativity between two bonded atoms is what determines the type of bond they form. Think of it as a sliding scale of sharing:

• Nonpolar Covalent Bond: If the electronegativity difference is small (like between two carbon atoms or carbon and hydrogen atoms), they share the electrons pretty equally. It’s a peaceful, harmonious bond.
• Polar Covalent Bond: If the electronegativity difference is moderate (like between oxygen and hydrogen), one atom pulls the electrons closer, but doesn’t completely steal them. This creates partial charges, which we will discuss later.
• Ionic Bond: If the electronegativity difference is huge (like between sodium and chlorine in table salt), one atom completely steals the electrons from the other. It’s a full-blown transfer, creating ions (charged atoms) and a strong electrostatic attraction. So, Sodium Chloride is a classic ionic compound.

Polar Covalent Bonds: Sharing is Caring, But Not Equally

• The Great Electron Heist (Kind Of): Okay, so imagine two atoms holding hands (bonded, if you want to get technical). They’re sharing electrons, which is the essence of a covalent bond. But what happens when one atom is just a tad bit greedier than the other? That’s where polar covalent bonds come in. It’s not a full-on snatch-and-grab like in ionic bonds, but more like a gentle tug-of-war where one atom pulls the electrons closer. This happens when there is a significant difference in electronegativity – but not so significant that you end up with ions!

• Introducing δ+ and δ-: The Partial Charge Crew: Because the electrons are spending more time with the greediest atom, that atom gets a slightly negative charge. We call it a partial negative charge, represented by the symbol δ- (delta minus). The other atom, feeling a bit electron-deprived, gets a partial positive charge (δ+). Think of it like sharing a pizza. If one person eats most of the slices, they’re not completely owning the pizza, but they’re definitely enjoying it more.

• Examples in the Wild: From Hydrochloric Acid to Sugars: Let’s look at some real-world examples. Hydrogen chloride (H-Cl) is a classic. Chlorine is much more electronegative than hydrogen, so the shared electrons hang out closer to the chlorine atom. This makes the chlorine end slightly negative (δ-) and the hydrogen end slightly positive (δ+). Carbon-oxygen bonds (C-O) are another common example, like in sugars. Oxygen hogs the electrons, making it partially negative and the carbon partially positive.

• Polar vs. Nonpolar: Knowing the Difference: So, how do you know if a bond is polar or not? It boils down to electronegativity differences. If the electronegativity difference is small (like between carbon and hydrogen, C-H), the bond is considered nonpolar. The electrons are shared relatively equally. But, if the difference is larger (like between oxygen and hydrogen, O-H), the bond is polar. One atom has a clear advantage in the electron tug-of-war. Polar bonds create partial charges, while nonpolar bonds keep things electrically balanced.

Dipole Moment: Putting a Number on Polarity – It’s More Than Just a Feeling!

Okay, so we know that polarity is about how unevenly electrons are shared. But how do we really know how polar something actually is? Enter the dipole moment, chemistry’s way of putting a number on the “pull” of electrons. Think of it as a tiny arrow pointing from the positive end of a molecule to the negative end. The longer the arrow, the more polar the molecule! It’s not just a feeling; it’s a measurable quantity that tells us exactly how much charge is separated and over what distance. It’s like giving polarity a report card, and the dipole moment is the grade!

Decoding the Dipole Moment Formula: μ = q * d

Ready for a little math? Don’t worry; it’s not as scary as balancing redox reactions! The dipole moment (represented by the Greek letter mu, μ) is calculated using a simple formula: μ = q * d.

• q stands for the magnitude of the partial charges (the δ+ and δ- we talked about before). The bigger the partial charges, the stronger the “pull,” and the larger the dipole moment. Think of it like this: A bigger difference in electronegativity results in a larger charge separation and a more powerful dipole.
• d stands for the distance between the charges. The further apart the positive and negative charges are, the larger the dipole moment. Imagine stretching that “arrow” further; the effect becomes more significant, right?

Molecular Geometry: Shape Matters, Like, A Lot!

Now, things get a bit trickier – and way more interesting! It’s not enough to simply have polar bonds to have a dipole moment. A molecule can have polar bonds, but because of its shape (or molecular geometry), those individual bond polarities can cancel each other out. It’s like a tug-of-war where everyone pulls with equal force in opposite directions – nobody moves!

• Bond Polarity is Key: A molecule needs to have polar bonds in the first place to even be considered for having a dipole moment. More polar bonds potentially mean a larger dipole moment – but geometry can still play spoiler.

• Symmetry Can Be a Killer: Consider carbon dioxide (CO₂). It has two polar carbon-oxygen bonds. But because the molecule is linear and perfectly symmetrical, the dipole moments of the two bonds cancel each other out. The net dipole moment of CO₂ is zero. It’s nonpolar, believe it or not!

  Now, think about water (H₂O). Like CO₂, it has polar bonds (oxygen is much more electronegative than hydrogen). But water is bent, not linear. Because of this, the individual bond dipoles don’t cancel. Instead, they add up to give water a significant net dipole moment, making it a very polar molecule! It’s the poster child for polarity!

Vector Addition: Adding Up the Arrows of Polarity

So, how do we figure out if a molecule’s polar overall? We use something called vector addition. Basically, we treat each bond dipole as an arrow and add them together like vectors in physics. If the arrows cancel each other out, the molecule is nonpolar. If they add up to a net arrow, the molecule is polar. It’s like a treasure hunt where following all the arrows leads you to the total “polarity treasure.” A fun way to think about it!

Water (H₂O): The Quintessential Polar Molecule

Alright, let’s talk about water—H₂O—the MVP of molecules! Ever wonder why water is so special? It all boils down (pun intended!) to its polarity. It’s not just some random arrangement of atoms; the molecule is built like a master architect planned it this way, specifically to become the world’s greatest solvent.

Geometry Matters: Bent But Not Broken!

The structure of water isn’t linear, like, say, carbon dioxide (O=C=O). Instead, it’s bent! The two hydrogen atoms aren’t directly opposite each other. This bend is critical. Now, oxygen is like the electron hog of the periodic table—super electronegative. This simply means it yanks the electrons from those poor hydrogens closer to itself.

Partial Charges: A Little Bit Positive, A Little Bit Negative

Because oxygen is hogging the electrons, it gets a slightly negative charge (δ-). This is because it has more electron density than it usually would. Consequently, the hydrogens end up with slightly positive charges (δ+). Imagine them as tiny, cheerful magnets! This unequal sharing of electrons is what makes water a polar molecule.

The High Dipole Moment: Water’s Superpower

Thanks to its bent shape and electronegativity differences, water has a high dipole moment. Remember, a dipole moment is a measure of how polar a molecule is. A high dipole moment means water is a polarity powerhouse! It’s like the Michael Jordan of molecules when it comes to attracting other charged particles. This is why it’s so good at dissolving stuff, especially ionic compounds (like salt) and other polar molecules.

Water: The Solvent of Life

This polarity is why water is such an amazing solvent, dissolving tons of things. This is especially crucial for biological molecules. Think about it: our bodies are mostly water! Water’s polarity helps transport nutrients, remove waste, and facilitate countless chemical reactions essential for life. Without water’s unique ability to dissolve and interact with biological molecules, life as we know it simply wouldn’t be possible. That’s the magic of water for you—a simple molecule with a powerful impact, all thanks to its polarity!

Hydrogen Bonding: The Ultimate Intermolecular Connection!

Alright, folks, let’s dive into the world of hydrogen bonding – the VIP of intermolecular forces! Forget regular ol’ attraction; this is where molecules get seriously cozy. We’re talking about a special interaction, almost like a molecular hug, that happens when a hydrogen atom, already attached to a super-electronegative atom like oxygen, nitrogen, or fluorine, gets drawn to a lone pair of electrons on another electronegative atom. Think of it as hydrogen reaching out for a little extra love!

So, what’s the recipe for this intermolecular romance? You need two key ingredients: a hydrogen bond donor and a hydrogen bond acceptor. The donor is your hydrogen atom already bonded to one of those electronegative powerhouses (O, N, or F). This makes the hydrogen partially positive (δ+), eager to mingle. The acceptor is another O, N, or F, sporting a lone pair of electrons just itching to form a connection. It’s like a molecular dating app – you need both parties to be present and have the right “chemistry”!

Now, let’s talk about the rockstar of hydrogen bonding: water! Water’s amazing properties, like its surprisingly high boiling point and its ability to support life, are almost entirely thanks to hydrogen bonds. Each water molecule can form hydrogen bonds with up to four other water molecules, creating a vast network of interactions. This is why water has such a high surface tension (think of water striders walking on water) and why it takes so much energy to boil (all those bonds need breaking!).

But water isn’t the only player in the hydrogen bonding game. Alcohols, those organic molecules with an -OH group, are also big fans. Like water, alcohols can both donate and accept hydrogen bonds. However, there’s a slight difference: while water can form a 3D network of hydrogen bonds, alcohols often have a nonpolar alkyl group attached, which can somewhat disrupt this network. This means that while alcohols can still form hydrogen bonds, their strength and impact on physical properties might be slightly different compared to water. Still, they’re close cousins in the hydrogen bonding family!

Alcohols (R-OH): Polar Cousins of Water

Ever wondered what makes your hand sanitizer effective or that beer so, well, *beer-like?* Let’s talk about alcohols! They’re like water’s slightly quirky cousins, sporting a hydroxyl group (-OH) that makes them surprisingly interesting on the molecular level.

The Hydroxyl Hookup: Why Alcohols are Polar

Think of an alcohol molecule as a hydrocarbon chain (that “R” in R-OH) playing host to a hydroxyl group (-OH). Now, oxygen is a bit of a greedy electron hog. It’s far more electronegative than both carbon and hydrogen. This electron-hogging behavior creates a polar bond within the O-H group. The oxygen gets a slightly negative charge (δ-), while the hydrogen gets a slightly positive charge (δ+). This uneven distribution of charge makes the alcohol molecule polar overall.

Hydrogen Bonding: Alcohols Gone Social

Because of that polar O-H bond, alcohols can engage in hydrogen bonding. They can bond with other alcohol molecules, creating little molecular cliques. They can also mingle with water molecules, making them fairly sociable. This hydrogen bonding has a massive impact on their properties.

• Solubility: Small alcohols like methanol and ethanol are completely miscible (mixable) with water. They’re basically the life of the party, easily forming hydrogen bonds with water molecules. However, as the carbon chain gets longer and bulkier, the alcohol becomes less soluble in water. It’s like inviting someone to a party who only speaks a different language – they just don’t mix as well!
• Boiling Points: Hydrogen bonding also leads to higher boiling points than you might expect for their size. It takes more energy to break those intermolecular hydrogen bonds and get the alcohol molecules to vaporize.

Alcohol Family Dynamics: A Polarity Comparison

Not all alcohols are created equal. Polarity and hydrogen bonding ability can vary depending on the structure:

• Methanol (CH3OH): This one is the simplest alcohol and highly polar. With only one carbon, it is completely miscible in water.
• Ethanol (CH3CH2OH): Still quite polar and miscible with water, ethanol is known as the alcohol in alcoholic beverages.
• Long-Chain Alcohols (e.g., Octanol, CH3(CH2)7OH): As the carbon chain gets longer, the nonpolar hydrocarbon portion dominates. These alcohols are much less polar and less soluble in water. They prefer hanging out with other nonpolar molecules.

Intermolecular Forces: The Glue That Holds Matter Together

Ever wondered why some things stick together and others don’t? It’s not just magic, it’s intermolecular forces (IMFs)! Think of them as the shy but powerful forces acting between molecules – the “glue” holding everything together. These forces can be attractive (bringing molecules closer) or repulsive (pushing them away), but we’re mostly interested in the attractive ones because they determine so much about how stuff behaves. Ready to dive in? Let’s unpack the different types of these molecular hugs!

Types of Intermolecular Forces

1. Dipole-Dipole Forces: Imagine two magnets coming together – that’s kinda like dipole-dipole forces! These happen between polar molecules, where one end is a bit positive and the other is a bit negative (remember those partial charges?). The positive end of one molecule is attracted to the negative end of another, like tiny little magnets aligning. It’s all about attraction, baby!

2. London Dispersion Forces (Van der Waals Forces): These are the universal IMFs, present in every single molecule, even the nonpolar ones. How? Electrons are always moving, and sometimes they bunch up on one side of a molecule, creating a temporary, fleeting dipole. This temporary dipole can then induce a dipole in a neighboring molecule, resulting in a weak, short-lived attraction. Think of it like this: even the wallflower at the party can get a dance if the music is right! These are especially important in nonpolar molecules, where they’re often the only IMFs present.

3. Hydrogen Bonding: The superhero of IMFs! This is a special, extra-strong type of dipole-dipole interaction that occurs when you have a hydrogen atom bonded to a super electronegative atom like oxygen (O), nitrogen (N), or fluorine (F). The partially positive hydrogen is then attracted to a lone pair of electrons on another O, N, or F. It’s like a super-charged magnet! Think water molecules clinging to each other like gossip buddies.

Polarity’s Role in Intermolecular Force Strength

So, where does polarity fit into all this? Well, it’s the VIP! The more polar a molecule is, the stronger its dipole-dipole forces will be. That’s because the partial charges are bigger, leading to a stronger attraction. Nonpolar molecules are stuck with the weaker London dispersion forces (although they can still add up!). Polar molecules generally experience stronger IMFs than nonpolar molecules, assuming they’re of similar size and shape. Size matters, too—bigger molecules have more electrons, leading to stronger London dispersion forces!

IMFs and Physical Properties: Boiling Points and Beyond

Okay, so we know what IMFs are, but why should we care? Because they have a huge impact on physical properties like boiling point and melting point! The stronger the IMFs, the more energy (heat) you need to overcome those attractions and separate the molecules. So, substances with strong IMFs have higher boiling points and melting points.

Think about it: water has strong hydrogen bonds, which is why it’s a liquid at room temperature and takes a lot of energy to boil. Methane (CH₄), on the other hand, is a nonpolar molecule with only weak London dispersion forces, so it’s a gas at room temperature and boils at a chilly -161°C. Understanding IMFs helps us predict and explain why different substances have such different properties. And that, my friends, is pretty cool!

Polarity and Solubility: “Like Dissolves Like” – A Matchmaking Game for Molecules

• Like dissolves like, ever heard that one? It’s the golden rule of solubility, a matchmaking principle for molecules. Think of it as chemistry’s version of finding your soulmate – polar molecules want to hang out with other polar molecules, and nonpolar molecules prefer their own kind. Imagine trying to mix oil and water; they just don’t get along, right? That’s polarity at play!

• Alcohols are interesting because they’re a bit like molecular chameleons. Small alcohols, like methanol and ethanol, love water. They’re totally miscible (mixable in all proportions) because they can form hydrogen bonds with water molecules. But as you add more carbon atoms, making the alcohol chain longer and more nonpolar, their love for water starts to fade. Suddenly, they’re more interested in hanging out with nonpolar solvents. It’s like they start developing a preference for the “other side.”

• So, what’s really going on here? It all boils down to intermolecular forces (IMFs). For a solute (the thing being dissolved) to dissolve in a solvent (the thing doing the dissolving), the IMFs between the solute and solvent must be strong enough to break apart the existing IMFs within each substance. Imagine trying to crash a party – you need to be charming enough to get past the bouncer (the IMFs holding the molecules together). If the solute-solvent interactions are weaker, the solute will stay cliqued up with its own kind, refusing to dissolve.

• Let’s bring this home with some real-world examples:

  • Salts in Water: Table salt (NaCl) is an ionic compound, which is extremely polar. That’s why it dissolves so well in water, another polar molecule. The positive sodium ions (Na+) and negative chloride ions (Cl-) are strongly attracted to the partially negative oxygen and partially positive hydrogen atoms in water, respectively. The water molecules surround the ions, breaking apart the crystal lattice and dispersing them throughout the solution.
  • Oils in Water: On the flip side, oils are mostly made up of long chains of carbon and hydrogen atoms, making them very nonpolar. When you try to mix oil and water, the water molecules are much more attracted to each other (through hydrogen bonding) than they are to the oil molecules. The water molecules squeeze the oil molecules out, resulting in two separate layers. It’s a classic case of “birds of a feather flock together.”
  • Sugars in Water: Sugars like sucrose (table sugar) have many –OH (hydroxyl) groups that allow them to form extensive hydrogen bonds with water, which allows them to dissolve well.

Polarity and Acidity/Basicity: It’s All About Those Electrons Feeling Good!

So, we’ve talked about how polarity dictates who likes to hang out with whom in the molecular world. But did you know it also plays a massive role in whether a molecule is an acid or a base? It’s true! And it all boils down to electron distribution and, frankly, molecular contentment. Think of it like this: molecules, just like us, want to be stable and happy. And for a conjugate base, that happiness often means having its negative charge nicely spread out and stabilized.

Now, how does polarity accomplish this magical feat? Well, imagine an electronegative atom – a real electron hog! – sitting near a negatively charged conjugate base. This electron-withdrawing neighbor can suck some of that negative charge density away, like a tiny vacuum cleaner for electrons. By spreading out the negative charge, it prevents any one atom from feeling too overwhelmed. The result? A more stable, less reactive conjugate base. And guess what? A stable conjugate base means a stronger acid! Basically, the easier it is for an acid to lose a proton (H+), the stronger the acid.

Inductive Effects: When Polarity Sends a Signal

This brings us to inductive effects, which are basically the molecular equivalent of gossip spreading through the neighborhood. If you have electron-withdrawing groups hanging around a molecule (like halogens or nitro groups), they’ll pull electron density towards themselves through the sigma bonds. This makes it easier for the molecule to donate a proton, boosting its acidity.

On the flip side, electron-donating groups (like alkyl groups) will push electron density towards the acidic part of the molecule. This makes it harder to lose that proton, decreasing acidity. It’s all about how those electrons are being shuffled around!

Acids, Bases, and Polar Personalities: A Few Familiar Faces

Let’s put some names to faces, shall we? Carboxylic acids (think vinegar – acetic acid) are classic examples of polar molecules acting as acids. The electronegative oxygen atoms in the carboxyl group help to stabilize the negative charge on the carboxylate ion after the acid loses a proton.

On the other hand, amines (like ammonia) are examples of polar molecules acting as bases. The nitrogen atom in an amine has a lone pair of electrons that it can use to accept a proton, making it a base. The polarity of the N-H bonds also contributes to its basicity.

Solvent Effects: Location, Location, Stability

Finally, let’s not forget the solvent, the medium in which the acidity/basicity drama unfolds. The polarity of the solvent can dramatically affect the acidity or basicity of a compound. Polar protic solvents (like water or alcohols) can stabilize ions through solvation, influencing the equilibrium of acid-base reactions. For instance, a polar solvent can better stabilize a charged conjugate base, promoting the acid’s dissociation. Conversely, in a nonpolar solvent, the stabilization of charged species is reduced, shifting the equilibrium. So, location matters, even for molecules!

Applications and Implications of Polarity

Polarity’s Starring Role in the Biological World

Ever wonder how your body manages to function with all those complex molecules floating around? Well, polarity is a major player!

• Protein Folding: A Molecular Hide-and-Seek. Think of proteins as intricate origami structures. Hydrophobic (water-fearing, or nonpolar) amino acids huddle together on the inside, away from water, while hydrophilic (water-loving, or polar) amino acids hang out on the surface, interacting with the watery environment. It’s like a molecular game of hide-and-seek, all driven by polarity! This drives the correct 3D shape to form for the protein to function properly.

• Enzyme-Substrate Interactions: A Polar Attraction. Enzymes are biological catalysts, and their ability to bind specifically to substrates (the molecules they act upon) often depends on polarity. Polar regions on the enzyme and substrate can attract each other, ensuring the right fit for the chemical reaction to occur. Imagine it as a lock and key system, but with electrical attraction in play!

• Cell Membrane Structure: A Polar Partnership. Our cells are surrounded by a membrane made of phospholipids. These molecules have a polar “head” and a nonpolar “tail”. This dual nature causes them to arrange themselves into a bilayer, with the polar heads facing the watery environment inside and outside the cell, and the nonpolar tails tucked away in the middle. Polarity is crucial for maintaining the membrane’s structure and controlling what enters and exits the cell. It’s like the ultimate bouncer at the cellular club.

Polarity: The Mastermind Behind Materials

Polarity isn’t just for biology; it’s also a design principle for the materials that shape our world.

• Designing Polymers: Polarity Powers. The properties of polymers (long chains of repeating units, like in plastics) are heavily influenced by the polarity of their monomers. Polarity can affect a polymer’s strength, flexibility, adhesion, and even its melting point. By carefully selecting monomers with specific polar characteristics, scientists can create polymers with tailored properties for various applications.

• Surfactants: The Soap Opera Stars. Surfactants are molecules with both polar and nonpolar regions. This unique structure allows them to reduce surface tension between liquids and stabilize emulsions (mixtures of liquids that don’t normally mix, like oil and water). Soap, for example, is a surfactant that uses its nonpolar region to grab onto grease and its polar region to dissolve in water, washing away the dirt.

Environmental Impacts: Polarity’s Downside

Polarity also plays a role in environmental issues, particularly concerning pollutants.

• Solubility of Pollutants in Water: A Polar Problem. Polar pollutants are more likely to dissolve in water, which can significantly impact their transport and fate in the environment. This means they can spread more easily through rivers, lakes, and groundwater, potentially contaminating drinking water sources and affecting aquatic ecosystems. It’s a reminder that even seemingly harmless substances can become environmental hazards due to their polarity.

So, next time you’re wondering why water is so good at dissolving stuff or why alcohols have such interesting properties, remember that little O-H bond and its polar nature. It’s a tiny thing that makes a big difference in the world around us!